GIFT   OF 
MICHAEL  REESE 


THE  NATUKE  OF 
SOLUTION 


BY 

HAEKY  C.  JONES 

LATE   PROFESSOR  OP  PHYSICAL  CHEMISTRY  IN  TBEE 
JOHNS  HOPKINS  UNIVERSITY 

WITH  A  BIOGRAPHICAL  SKETCH  BY 

E.  EMMET  REID 

PROFESSOR  OF  ORGANIC  CHEMISTRY  IN  THB 
JOHNS  HOPKINS   UNIVERSITY 

AND  TRIBUTES  BY 

PROFESSORS  ARRHENIUS, 
OSTWALD  AND  WOODWARD 


ILLUSTRATED 


NEW  YORK 

D.  VAN  NOSTRAND  COMPANY 

25  PARK  PLACE 

1917 


COPYRIGHT,    1917,  BY 
D.   VAN    NOSTRAND    COMPANY 


THE  •   PLIMPTON  •  PRESS 
NORWOOD-MASS-U-S-A 


PREFACE 

No  subject  in  chemistry  has  received  more  attention, 
especially  during  the  last  quarter  of  a  century,  than  that 
of  solution.  This  is  due  primarily  to  the  fundamental 
significance  of  solution  for  chemical  science. 

Solutions  in  the  broad  sense  of  the  term  are  funda- 
mental not  only  for  chemistry,  but  for  geology  and  the 
various  branches  of  biology.  Matter  hi  the  pure,  homo- 
geneous condition  is  relatively  inert.  It  becomes  active 
when  mixed  in  a  certain  way  with  other  matter  hi  the 
same  or  in  a  different  state  of  aggregation  —  when  dis- 
solved. 

Since  solution  is  so  fundamental  for  the  natural  sciences 
in  general,  and  for  chemistry  in  particular,  we  must  know 
what  solutions  are,  if  we  would  ever  make  these  various 
branches  of  science  exact.  Since  chemistry  has  to  do 
largely  with  the  science  of  solution,  it  can  become  an  exact 
science  only  by  the  science  of  solution  becoming  exact. 
We  must  first  know  what  is  the  real  condition  of  matter 
in  solution.  What  laws  does  it  obey?  Is  Jie  dissolved 
substance  combined  with  the  solvent,  and  if  so  with  how 
much  of  it? 

As  we  shall  see,  many  of  these  questions  have  now  been 
answered  and  the  relations  between  solutions  and  £•  ses 
accurately  established.  This  is  of  the  greatest  im^or- 
tance.  We  really  know  something  about  matter  in  the 
gaseous  state,  and  we  can  now  apply  this  knowledge  to 
matter  hi  the  dissolved  condition;  and  this  has  done 
more  than  any  other  one  discovery  to  place  the  science 
of  solution  upon  an  exact  basis. 


376133 


iv  PREFACE 

The  very  intensive  work  done  in  this  laboratory  on  the 
nature  of  solution  has  been  carried  out  largely  with  the 
aid  of  grants  from  the  Carnegie  Institution  of  Washington, 
without  which  much  of  the  work  could  not  have  been 
completed  and  many  of  the  results  obtained  could  not 
have  been  published. 

The  Carnegie  Institution  has  generously  allowed  me  to 
make  free  use  of  Chapter  VII,  No.  210,  Publication  of  the 
Carnegie  Institution  of  Washington,  especially  in  preparing 
chapters  XI,  XIV,  and  XV  of  this  work.  The  Franklin 
Institute  of  Philadelphia  has  also  allowed  me  to  use  parts 
of  one  of  my  recent  papers  hi  their  Journal. 

In  preparing  Chapter  II  many  references  to  the  liter- 
ature have  been  obtained  with  the  aid  of  the  admirable 
little  book  by  Walden,  "Die  Losungstheorien  in  ihrer 
geschichtlichen  Aufeinanderfolge." 

In  addition  to  true  solutions  we  have  colloidal  solu- 
tions, which  have  come  very  prominently  to  the  front 
in  the  last  ten  or  fifteen  years.  While  most  of  our 
knowledge  in  this  field  is  still  empirical,  yet  much  of 
value  has  been  learned,  and  the  importance  of  colloids 
for  the  science  of  living,  as  well  as  of  dead  matter, 
is  beginning  to  be  realized.  In  preparing  the  chapter 
on  colloids,  I  would  express  my  indebtedness  to  the 
comprehensive  work  of  Freundlich,  "Kapillarchemie," 
especially  for  references  to  the  literature  on  col- 
loidal chemistry.  This  is  the  great  hand-book  in  this 
field. 

Many  references  to  the  literature  on  solutions  have  been 
obtained  with  the  aid  of  my  "  Elements  of  Physical  Chem- 
istry/' 4th  Ed.  (The  Macmillan  Co.,  N.Y.),  to  which  the 
student  desiring  a  systematic  discussion  of  physical  chem- 
istry is  referred. 

The  present  work  is  not  a  text-book,  but  a  general 
discussion  of  some  of  the  more  important  properties  of 
solutions,  true  and  colloidal.  It  is  therefore  written  in  a 
nonmathematical,  indeed,  largely  in  a  semi-popular  style. 


PREFACE  V 

It  is  hoped  that  this  work  may  interest  students  of  the 
various  branches  of  science  to  go  on  into  the  real  physical 
chemistry  of  solutions,  and  from  this  into  physical  chem- 
istry hi  its  broadest  sense. 

HARRY  C.  JONES. 


BIOGRAPHICAL  SKETCH 

"  The  man  or  woman,  who  does  work  worth  doing,  is  the  man  or  woman  who  livea, 
breathes  and  sleeps  that  work;  with  whom  it  is  ever  present  in  his,  or  her 
soul;  whose  ambition  it  is  to  do  it  well  and  feel  rewarded  by  the  thought  of 
having  done  it  well.  That  man,  that  woman,  puts  the  whole  country  under 
an  obligation."  —  RUSKIN. 

FEW  men  of  science  have  ever  worked  with  such  consuming  intensity 
as  did  the  late  Professor  Harry  Clary  Jones.  His  tremendous  energy 
was  at  once  an  inspiration  and  a  despair  to  those  associated  with  him. 
His  enthusiasm  was  contagious,  yet  few  could  hold  the  pace  he  set. 

In  the  fifty  years  of  his  life  he  lived  more  than  most  who  reach  three 
score  and  ten,  and  into  his  twenty  years  of  scientific  activity  he  crowded 
enough  work  to  last  several  men  a  lifetime.  Work  was  his  vocation, 
his  vacation,  his  duty,  his  dissipation,  his  life,  his  death. 

He  was  a  large  man  possessing  great  physical  force  and  a  strength 
which  never  seemed  to  tire.  He  worked  long  hours  at  his  laboratory 
and  went  home  to  read  proof.  In  summer  he  would  go  away  for  a  vaca- 
tion, but  would  spend  it  writing  a  book;  when  a  bright  Saturday  after- 
noon came,  he  would  get  away  to  the  country,  but  to  spend  the  hours 
riding  over  his  three  farms  telling  his  farmers  how  to  raise  more  corn  and 
wheat  on  his  fertile  fields. 

There  is  something  strange  about  this  farmer  boy  leaving  the  well- 
tilled  acres  of  his  father  and  his  father's  father  and  coming  to  the  city, 
to  the  University,  where  by  unremitting  labor  he  forged  his  way  to  the 
front.  In  sight  of  his  birth-place,  a  stately  brick  house  built  by  his 
grandfather,  still  stands  a  tiny  country  school-house  where  his  early 
education  was  obtained,  and,  hard  by,  a  little  Methodist  church  of  which 
he  was  a  member  hi  his  youth,  and  to  which  he  always  contributed.  His 
remains  repose  hi  this  church-yard  hi  sight  of  the  school-house  and  the 
farm.  Well  nigh  the  entire  population  of  the  surrounding  country 
came  together  to  do  him  last  honors.  They  had  heard  of  his  scientific 
achievements  but  to  them  he  was  a  fellow  fanner,  a  leading  citizen,  and 
a  benefactor  of  the  community. 

To  the  end  of  his  life  he  managed  that  farm  efficiently  and  profitably, 
and  to  it  added  two  adjoining  farms  which  he  managed  equally  well,  so 
that  he  really  was  one  of  the  large  farmers  hi  one  of  the  most  fertile 
portions  of  Maryland,  but  farming  was  never  allowed  to  interfere  with 
his  scientific  work. 


viii  BIOGRAPHICAL  SKETCH 

Professor  Jones  was  born  near  New  London,  a  village  in  Frederick 
County,  Maryland  (U.  S.  A.),  November  llth,  1865,  and  there  spent 
his  early  years.  The  reading  of  one  of  TyndalTs  books  on  science 
awakened  his  interest  in  physical  science  and  gave  him  the  impulse  to 
seek  a  college  education.  At  the  age  of  twenty  he  entered  Johns  Hop- 
kins University  as  a  "special  student,"  his  preparation  not  being  suffi- 
cient for  regular  matriculation,  but  his  deficiencies  were  soon  made  up 
and  he  passed  to  the  degree  of  bachelor  of  arts  hi  1889.  His  untiring 
industry  and  mental  alertness  were  already  conspicuous  hi  his  under- 
graduate years. 

He  remained  at  the  University  for  graduate  work  taking  chemistry 
as  major  with  physics  and  mathematics  as  minor  subjects. 

Two  scholarships  as  an  undergraduate,  and  two  more  scholarships 
and  a  fellowship  hi  his  three  years  as  a  graduate  student  testify  to  his 
merits.  His  first  research  was  carried  out  under  Professor  H.  N.  Morse, 
"On  the  Action  of  Metallic  Cadmium  on  the  Halogen  Salts  of  Cadmium" 
and  on  "A  Redetermination  of  the  Atomic  Weight  of  Cadmium." 

He  received  the  degree  of  Doctor  of  Philosophy  from  the  Johns  Hop- 
kins University  hi  June,  1892.  He  felt  drawn  to  physical  chemistry, 
which  was  at  that  time  becoming  differentiated  from  the  older  chemistry 
and  which  had  come  into  prominence  on  account  of  the  work  of  van't 
Hoff,  Ostwald,  and  Arrhenius.  Hence  Dr.  Jones  crossed  the  ocean  to 
study  with  these  three  leaders,  and  worked  with  Ostwald  at  Leipzig, 
Arrhenius  at  Stockholm,  and  van't  Hoff  at  Amsterdam.  The  two  years 
spent  in  this  way  made  him  a  thorough-going  physical  chemist,  and  into 
physical  chemistry  he  flung  himself  with  consiuning  ardor.  The  reader 
will  find  from  the  contents  of  this  book  how  highly  Dr.  Jones  valued  these 
three  men  and  how  their  ideas  permeated  his  thinking.  He  renewed 
his  acquaintance  with  them  in  his  frequent  summer  trips  to  Europe. 
Professor  Ostwald  and  Professor  Arrhenius  were  his  close  friends  till 
his  death. 

In  the  fall  of  1894  he  returned  to  the  Johns  Hopkins  University  as  a 
Fellow  by  Courtesy.  This  was  a  purely  honorary  position  giving  him  the 
use  of  the  laboratory  but  requiring  nothing  of  him.  He  gave  a  few 
lectures  hi  physical  chemistry  but  spent  the  most  of  his  time  on  research. 
The  next  year  he  was  appointed  Instructor  hi  Physical  Chemistry,  being 
promoted  hi  1898  to  Associate.  In  1900  he  was  made  Associate  Profes- 
sor and  hi  1903,  Professor  of  Physical  Chemistry  in  Johns  Hopkins  Uni- 
versity, which  position  he  retained  to  the  tune  of  his  death.  In  1902  he 
married  Miss  Harriet  Brooks,  a  member  of  one  of  the  old  families  of 
Baltimore,  who  survives  him. 

He  carried  out  a  piece  of  research  in  the  laboratory  of  Arrhenius  on  the 
hydrates  of  sulphuric  acid,  mention  of  which  is  made  on  pages  46-48  of 
the  present  volume.  Thus  began  his  work  on  hydrates  hi  solution.  After 


BIOGRAPHICAL  SKETCH  ix 

his  return  to  this  country  he  began  a  systematic  study  of  the  deviations, 
found  in  more  concentrated  solutions,  from  the  laws  of  gas  pressure 
and  the  theory  of  electrolytic  dissociation.  From  1903  on  he  was  en- 
abled to  enlarge  this  work  by  generous  grants  from  the  Carnegie  Institu- 
tion of  Washington.  These  grants  were  continued  till  the  time  of  his 
death  and  the  same  Institution  has,  since  his  death,  provided  for  the 
rounding  out  and  completion  of  several  of  his  investigations. 

His  earlier  work  was  on  the  abnormalities  hi  the  molecular  depression 
of  the  freezing  point  in  aqueous  solutions  of  certain  substances  which 
separate  from  solution  with  water  of  crystallization.  In  1900  he  offered 
the  folio  whig  as  a  tentative  explanation:  "In  concentrated  solutions 
these  substances  must  take  up  a  part  of  the  water  forming  complex 
compounds  with  it  and  thus  removing  it  from  the  field  of  action  as  far 
as  freezing-pomt  lowering  is  concerned.  .  .  .  The  lowering  of  the  freezing 
point  is  thus  abnormally  great,  because  a  part  of  the  water  is  no  longer 
present  as  solvent.  .  .  .  We  do  not  put  forward  the  above  suggestion 
as  the  final  statement  of  a  theory,  but  only  as  a  tentative  explanation 
which,  however,  seems  to  account  for  the  experimental  facts  which  have 
been  brought  to  light."  This  was  his  first  statement  of  the  hydrate 
theory  which  was  the  central  thought  hi  all  of  his  subsequent  work. 
His  work  spread  out  in  ever-widening  circles,  but  the  idea  of  hydrates, 
or  solvates,  to  use  the  more  general  term,  always  remained  the  center. 

In  1901  he  undertook  the  study  of  the  conductivity  and  dissociation 
of  electrolytes  and  their  temperature  coefficients.  This  study  was  car- 
ried over  into  non-aqueous  and  mixed  solvents,  and  parallel  investiga- 
tions were  carried  out  on  the  viscosity  of  the  solutions,  showing  the 
important  bearing  of  viscosity  on  conductance  and  arriving  at  an  ex- 
planation of  various  phenomena  connected  with  viscosity,  e.g.  negative 
viscosity  coefficients.  Using  the  conductivity  method,  many  inorganic 
substances  and  organic  acids  were  studied  in  water  solution.  By  a  vast 
number  of  measurements  it  was  shown  that  these  temperature  coeffi- 
cients are  greater  in  the  case  of  the  more  hydrated  substances  indicating 
a  breaking  down  of  the  hydrates  with  elevation  of  temperature. 

A  large  amount  of  work,  extending  over  several  years,  was  done  on  the 
absorption  spectra  of  solutions.  This  involved  observations  on  thousands 
of  solutions  hi  several  solvents  and  in  many  mixed  solvents.  Four  of 
the  ten  monographs  published  by  the  Carnegie  Institution  of  Washing- 
ton, embodying  his  investigations,  are  devoted  to  this  spectrographic 
work.  More  recently  this  was  extended  farther  down  the  spectrum, 
as  well  as  made  more  quantitative  in  character,  by  the  use  of  a  radiomi- 
crometer  with  a  fine  grating  spectroscope.  With  this  apparatus  the 
ionization  constants  of  indicators  and  the  light  absorption  coefficients 
of  solutions  were  studied  in  addition  to  the  phenomena  of  solvation. 
The  spectrographic  evidence  showed  that  in  general,  the  light  absorp- 


x  BIOGRAPHICAL  SKETCH 

tion  is  characteristic  of  the  combination  of  the  solute,  i.e.  the  solvate 
with  the  solvent  rather  than  of  the  solute  alone. 

All  hi  all,  he  investigated  sixteen  lines  of  evidence  bearing  on  the  pres- 
ence of  solvates  hi  solution.  In  this  he  was  aided  by  a  large  number 
of  advanced  students  and  assistants.  Counting  up  the  tune  his  various 
coworkers  put  on  this  study  would  make  one  hundred  and  fifteen  years 
for  one  worker.  A  summary  of  all  these  investigations  is  given  hi  Chapter 
VII,  Publication  210,  of  the  Carnegie  Institution  of  Washington.  The  last 
two  chapters  of  the  present  volume  are  mainly  an  exposition  of  his  own 
work  on  solvation.  His  own  conclusion  is:  "The  theory  of  electrolytic 
dissociation,  supplemented  by  the  theory  of  solvation,  is  then  not  simply 
a  theory  of  dilute  or  ideal  solutions,  but  a  theory  of  solutions  hi  general." 

Before  his  work  began,  the  idea  of  solvation  had  been  frequently 
suggested  and,  contemporaneously  with  his  work,  various  important 
investigations  have  been  carried  out  along  somewhat  similar  lines  by 
other  investigators,  yet  the  amount  and  variety  of  evidence  which  he 
obtained  on  the  existence  and  behavior  of  solvates  give  him  a  preeminent 
place  in  connection  with  the  theory  of  solvation.  His  work  has  gone  far 
toward  obtaining  general  recognition  of  the  great  importance  of  solva- 
tion in  many  phenomena,  not  only  in  the  field  of  chemistry,  but  also  of 
biology  and  other  sciences. 

He  was  a  man  of  marked  personality  with  strong  likes  and  dislikes. 
He  was  a  pioneer  and  a  promoter  of  physical  chemistry  hi  America,  a 
champion  of  a  cause,  an  advocate  rather  than  a  judge.  He  was  not 
easily  shaken  from  an  opinion  but,  on  sufficient  evidence,  would  quickly 
and  frankly  acknowledge  an  error. 

He  was  an  enthusiastic  teacher  and  a  remarkably  clear  lecturer;  he 
spoke  rapidly  but  the  right  word  seldom  failed  him.  Those  who  worked 
under  him  enjoyed  considerable  liberty  of  thought  and  action.  In  every 
way  he  showed  a  keen  interest  hi  them  and  was  quick  to  aid  them,  looking 
out  for  their  advancement  even  years  after  they  had  left  the  University. 

His  laboratory  was  a  place  of  intense  chemical  activity.  In  twenty 
years  fifty-seven  persons  carried  on  research  under  his  guidance,  about 
fifty  of  whom  did  their  doctorial  dissertation  work  with  him.  Many  of 
these  remained  one  or  two  years  as  Carnegie  assistants,  after  receiving 
the  doctor's  degree.  Professor  Jones  and  his  collaborators  published 
one  hundred  and  forty  articles  in  American,  German,  English,  and 
French  journals.  Ten  monographs  written  by  Professor  Jones,  embody- 
ing his  investigations,  have  been  published  by  the  Carnegie  Institution  of 
Washington  and  one  is  now  being  prepared. 

Professor  Jones  was  a  writer  of  books,  the  present  volume  being  the 
twelfth,  for  the  student  of  chemistry  and  the  general  reader  of  science. 
He  wrote  with  great  facility  and  amazing  rapidity,  in  one  case  writing  a 
book  in  six  weeks,  yet  his  style  was  good,  always  readable  and  clear. 


BIOGRAPHICAL  SKETCH  xi 

He  maintained  that  by  writing  slowly  one  loses  force  and  charm.  Many 
readers  have  found  his  more  popular  books  as  difficult  as  a  novel  to  lay 
aside.  Several  of  his  books  have  been  quite  successful,  his  "  Elements  of 
Physical  Chemistry"  having  gone  through  four  editions  and  having  been 
translated  into  Italian  and  Russian,  while  others  have  been  reprinted  sev- 
eral times.  All  hi  all  about  20,000  copies  of  Professor  Jones'  books  have 
been  sold.  ••••*  ~*i''$ 

A  bibliography  of  his  articles  and  books  is  to  be  found  hi  the  appendix 
to  this  volume.  The  most  of  this  is  taken  from  the  Johns  Hopkins 
University  Circular  of  February,  1916.  •• 

In  his  youth  he  studied  the  piano  and  debated  long  between  it  and 
science  but  never  touched  it  after  he  made  his  decision  for  science.  He 
retained  his  fondness  for  music  and  was  a  frequent  and  appreciative 
hearer  of  the  best.  He  was  a  genial  host  and  loved  to  entertain  his  friends 
in  his  well-appointed  home.  His  students,  hi  particular,  look  back  with 
pleasure  to  informal  dinners  with  him  and  Mrs.  Jones.  He  traveled 
extensively  hi  Europe  and  hi  America.  i-  i 

He  was  a  member  of  the  American  Chemical  Society,  the  American 
Physical  Society,  the  Franklin  Institute,  American  Philosophical  Society, 
and  Washington  Academy  of  Science,  and  honorary  member  of  the 
Brooklyn  Institute  of  Arts  and  Sciences. 

He  was  an  associate  editor  of  the  Zeitschrift  fur  physikalishe  Chemie, 
the  Journal  de  Chemie  physique  and  Journal  of  the  Franklin  Institute. 
In  1913  he  was  the  Longstreth  Medalist  of  the  Franklin  Institute.  His 
unremitting  work  and  an  inherited  tendency  to  nervousness  brought  on 
insomnia  and  melancholia  which  made  his  last  months  almost  unbearable 
and  led  to  his  untimely  death  on  April  9,  1916.  He  learned  many  things 
but  never  learned  to  rest. 

The  present  volume  was  written  by  Professor  Jones  during  the  last 
summer  of  his  life  and  put  into  the  hands  of  the  printer,  but  later  he 
withdrew  it  from  publication.  After  his  death  it  was  decided  to  issue 
it  as  a  memorial  volume.  The  writer  of  this  sketch  has  supervised 
bringing  it  out.  Professor  E.  C.  Bingham,  of  Lafayette  College,  has 
kindly  read  the  proof  and  made  valuable  suggestions.  The  burden  of 
correcting  the  proof,  verifying  references,  making  indexes,  etc.,  has  been 
borne  by  Dr.  Jones'  Carnegie  assistants,  Dr.  P.  B.  Davis  and  Dr.  H.  H. 
Lloyd,  and  thanks  are  due  them  for  then-  faithful  work. 

His  work  is  all  on  record  and  his  place  in  Chemistry  will  best  be 
decided  by  the  chemists  of  the  future,  but  a  few  tributes  from  some 
who  are  well  able  to  judge  are  not  out  of  place  at  this  time. 

E.  EMMET  REID. 
JOHNS  HOPKINS  UNIVBBSITT, 

BALTIMORE,  MD. 
January,  1917. 


CONTENTS 


CHAPTER  I 
IMPORTANCE  OP  SOLUTION  PAGE 

Importance  of  Solution  Early  Recognized 1 

Types  of  Solutions 1 

State  of  Aggregation  a  Function  of  Conditions 3 

Aqueous  Solutions  the  More  Important 4 

Present  Use  of  the  Term  Solution 5 

Solution  and  Chemical  Transformation 6 

Solution  Fundamental  for  all  Branches  of  Chemistry 7 

Dry  Substances  do  not  Generally  React  Chemically 7 

Dry  Chlorine  and  Dry  Sodium 8 

Dry  Hydrochloric  Acid  Gas  on  Dry  Carbonates 8 

Dry  Acids  on  Dry  Litmus 9 

Dry  Ammonium  Chloride  Sublimes  Undecomposed 9 

Dry  Hydrochloric  Acid  Gas  on  Dry  Ammonia  Gas • 

Dry  Sulphuric  Acid  on  Dry  Metallic  Sodium 11 

Importance  of  Solution  for  Physics 12 

Solution  and  Geology 13 

Solution  and  Biology 14 

Water  a  Remarkable  Compound 16 

Heat  Evolved  in  the  Formation  of  Water 17 

Physical  Properties  of  Water 17 

Importance  of  Solution  Justifies  the  Work  Expended  upon  It 19 

CHAPTER  H 
EARLIER  VIEWS  AS  TO  THE  NATURE  OP  SOLUTION 

Sir  Isaac  Newton 20 

Boerhaave 21 

Wallerius 22 

Lavoisier 22 

Fourcroy 23 

Klaproth 23 

BerthoUet 23 

Thomson 25 

Grotthuss 25 

Berzelius'  Electrochemical  Theory 27 

Berzelius'  Theory  of  Solution 29 

Gay-Lussac 30 


xvi  CONTENTS 

Williamson 31 

Clausius 33 

Kopp 35 

Guldberg  and  Waage 35 

Valson 36 

Favre  and  Valson 38 

Landolt 39 

Gladstone / 39 

Kohlrausch 40 

Berthelot 40 

Thomsen t 42 

Raoult. . .   : 43 

Mendeleeff 44 

Testing  Mendeleeff's  Hypothesis 46 

End  of  the  Qualitative  Period 48 

CHAPTER  III 
THE  OSMOTIC  PRESSURE  OF  SOLUTIONS 

Traube's  Method  of  Measuring  Osmotic  Pressure 50 

Pfeffer's  Preparation  of  Semi-permeable  Membranes 50 

Pfeffer's  Measurements 51 

Pfeffer's  Results 52 

Where  Pfeffer  Left  the  Subject  of  Osmotic  Pressure 53 

Other  Measurements  of  Osmotic  Pressure 55 

Work  of  Berkeley  and  Hartley  on  the  Osmotic  Pressures  of  Concentrated 

Solutions 55 

Results  Obtained 56 

Pfeffer's  Membranes  would  not  Withstand  Great  Pressures 58 

Electrical  Method  of  Preparing  Semi-permeable  Membranes 58 

Measurements  of  Osmotic  Pressure  made  by  Morse,  Frazer,  Holland  and 

Co-workers 59 

Results  with  Cane  Sugar 60 

Ratios  between  Gas  Pressure  and  Osmotic  Pressure 60 

Summary  of  Relations 62 

Results  with  Glucose 63 

Results  with  Mannite 64 

Osmotic  Pressure  of  an  Electrolyte 64 

Results  with  Lithium  Chloride 65 

Durability  of  the  Morse  Cells 66 

Cell  for  High  Osmotic  Pressures 66 

De  Vries'  Method  of  Measuring  the  Relative  Osmotic  Pressures  of  Solu- 
tions   67 

Method  of  Procedure 69 

Isotonic  Coefficients 69 

Limitations  of  the  Method 70 

Other  Methods  of  Determining  the  Relative  Osmotic  Pressures  of  Solu- 
tions. .  70 


CONTENTS  xvii 


CHAPTER  IV 

RELATIONS  BETWEEN  SOLUTIONS  AND  GASES  DEMONSTRATED  BT  VAN'T 

HOPE 

Law  of  Boyle  Applies  to  the  Osmotic  Pressures  of  Dilute  Solutions 73 

Law  of  Gay-Lussac  Applies  to  the  Osmotic  Pressure  of  Dilute  Solutions.  74 

Work  of  Soret 75 

Equality  of  Gas  Pressure  and  of  Osmotic  Pressure . 76 

Importance  of  the  Applicability  of  the  Gas  Laws  to  Solutions 78 

Exceptions  to  the  Above  Relations 79 


CHAPTER  V 

THE  THEORY  OF  ELECTROLYTIC  DISSOCIATION  AS  ANNOUNCED  BY 

ARRHENITJS 

Arrhenius  and  the  Dissociation  Theory 82 

Arrhenius  Points  out  Methods  of  Measuring  the  Magnitude  of  Dissocia- 
tion    83 

Dissociation  from  Freezing-point  Lowering  Compared  with  Dissocia- 
tion from  Conductivity 84 

Bearing  of  the  Theory  of  Electrolytic  Dissociation  on  Chemistry 86 

Opposition  to  the  Theory  of  Electrolytic  Dissociation 87 

Relations  between  Solutions  and  Gases  and  the  Theory  of  Electrolytic 

Dissociation  hold  only  for  Dilute  Solutions 88 


CHAPTER  VI 
DIFFUSION  IN  SOLUTION 

The  Phenomenon  of  Diffusion v. . .  ; . ; 90 

Diffusion  Caused  by  Some  Force 90 

Work  of  Graham  on  Diffusion 91 

The  Generalization  of  Fick 93 

Effect  of  Mass  on  Diffusion 93 

Temperature  Coefficients  of  Diffusion  —  Principle  of  Soret 96 

Osmotic  Pressure  Produces  Diffusion 96 

Importance  of  Osmotic  Pressure  for  Chemistry 99 

Bearing  of  Osmotic  Pressure  on  Biological  Phenomena 99 

Osmotic  Pressure  and  Geology 101 

CHAPTER  VH 

DEPRESSION  OF  THE  VAPOR-TENSION  OF  A  SOLVENT  BY  SUBSTANCES 
DISSOLVED  IN  IT 

Work  of  Faraday 103 

Investigations  of  Wullner 104 

Study  of  Vapor-tension  by  Walker. 104 


xviii  CONTENTS 

Raoult  and  the  Vapor-tensions  of  Solvents  and  of  Solutions 106 

Raoult's  Results  —  Effect  of  Concentration 107 

Effect  of  Temperature 108 

Effect  of  Nature  of  Dissolved  Substance 109 

Raoult's  Law  —  Effect  of  Nature  of  Solvent 110 

Frazer-Lovelace  Method  of  Measuring  Vapor-tension Ill 

The  Boiling-point  Method 114 

Boiling-point  Method  of  Beckmann 115 

Boiling-point  Apparatus  of  Jones 116 

Molecular  Weights  from  Boiling-point  Determinations 117 

Electrolytic  Dissociation  Measured  by  the  Boiling-point  Method 118 

Molecular  Weights  of  the  Metals  in  Mercury 119 


CHAPTER  VIII 
DEPRESSION  OF  THE  FREEZING-POINT  OP  A  SOLVENT  BY  THE  SOLUTE 

Investigations  of  Raoult 122 

Results  obtained  by  Raoult 123 

Raoult's  Law  of  Freezing-point  Lowering 126 

The  Freezing-point  Method  of  Beckmann 127 

Results  of  Molecular  Weight   Determinations  by  the   Freezing-point 

Method 128 

Electrolytic  Dissociation  Measured  by  the  Freezing-point  Method 129 

Improved  Freezing-point  Method 130 

Results  of  Measurements  of  Dissociation  by  the  Freezing-point  Method  130 

The  Three  Fundamental  Properties  of  Solutions 132 

Freezing  of  Saturated  Solutions 133 


CHAPTER  IX 

AQUEOUS  SOLUTIONS  OP  ACIDS,  BASES,  AND  SALTS  —  ELECTROLYTES 

What  is  an  Acid? 136 

What  is  a  Base? 137 

Neutralization  of  Acids  by  Bases 137 

Importance  of  this  Fact  for  Chemistry 139 

Heat  of  Neutralization  of  Acids  and  Bases 139 

Exceptions  Presented  by  Weak  Acids  and  Weak  Bases 141 

Theory  of  Electrolytic  Dissociation  as  a  Correlator  of  Facts 141 

Law  of  the  Thermoneutrality  of  Salts 142 

Importance  of  Energy  Changes  for  Chemistry 143 

Color  in  Solution 144 

Absorption  of  Light  Due  to  Resonance 145 

Color  in  Solution  May  Be  Ionic  or  Molecular 145 

Cause  of  Color  in  Solution. .  147 


CONTENTS  xix 

Color  Changes  and  Volumetric  Analysis  —  Indicators 148 

Stieglitz's  Views  in  Reference  to  Phenolphthalein 150 

Amphoteric  Electrolytes  —  Importance  in  Living  Organisms 151 

Amphotenism  and  Biological  Processes 152 

Hydrolysis  at  Ordinary  and  at  High  Temperatures 153 

Effect  of  Temperature  on  Hydrolytic  Dissociation 154 


CHAPTER  X 
SOME  ELECTRICAL  PROPERTIES  OF  AQUEOUS  SOLUTIONS  OP  ELECTROLYTES 

Faraday's  First  Law ..;..,, ...  - .-'. . . . .  V. .  „-. ..  .\ ....  156 

Faraday's  Second  Law 157 

Second  Law  of  Faraday  and  Chemical  Valence 157 

Nature  of  Chemical  Valence 158 

Experiment  Demonstrating  the  Nature  of  Chemical  Valence 159 

Electrolytes  Conduct  Only  by  Undergoing  Electrolysis 161 

The  Laws  of  Faraday  Apparently  Rigid  Laws  of  Nature 163 

The  Nature  of  Electrolysis 164 

Older  Theory 164 

The  Decomposition  Values  of  the  Ions 165 

Newer  Theory  of  Electrolysis 166 

Newer  Theory  and  Decomposition  Values 166 

Electrolysis  of  Water  a  Direct  Decomposition  by  the  Current 170 

Property  of  Solutions  of  Electrolytes  to  Conduct  the  Electric  Current. .  170 

The  Kohlrausch  Method  of  Measuring  Electrical  Conductivity 172 

Applying  the  Kohlrausch  Conductivity  Method 173 

Purification  of  Water 174 

Regulation  of  Temperature 175 

Some  Results  of  Conductivity  Measurements 176 

Law  of  Kohlrausch  and  Ostwald 177 

Dissociation  of  Electrolytes  Measured  by  the  Conductivity  Method 178 

Relation  between  Dissociation  and  Dilution 179 

Conductivity  and  Dissociation  at  Elevated  Temperatures 181 

How  Electrolytes  are  Dissociated  by  Water 182 

Ways  in  which  Ions  are  Formed 186 

Velocities  with  which  the  Ions  Move 188 

Apparatus  Designed  by  the  Author 188 

Effect  of  Concentration  and  Temperature 190 

Results  of  Measurements  of  the  Relative  Velocities  of  the  Elementary 

Ions 191 

Interpretation  of  the  Results 192 

Absolute  Velocities  of  the  Ions 194 

Velocities  of  Ions  and  of  Gaseous  Molecules 195 

Electromotive  Force  of  Primary  Cells 196 

Solution-tension  of  the  Metals 197 

Proof  of  the  Existence  of  Solution-Tension 198 

Values  of  the  Solution-tension  of  Certain  Metals 198 


XX  CONTENTS 

CHAPTER  XI 
SOLUTION  IN  NONAQUEOUS  AND  IN  MIXED  SOLVENTS 

Solvent  Power  of  Liquids 202 

Dissociating  Powers  of  the  Different  Solvents  —  Organic  Solvents 202 

Inorganic  Solvents 205 

Abnormal  Electrolytes 206 

Dissociating  Power  of  Solvents  and  Their  Dielectric  Constants 209 

Determination  of  the  Dielectric  Constants  of  Media 210 

Relation  between  the  Dissociating  Power  of  Solvents  and  their  own  Asso- 
ciation    211 

Effect  of  Rise  in  Temperature  on  the  Association  of  Liquids  and  on  Their 

Dissociating  Power 212 

Results  in  Mixed  Solvents 212 

Mixtures  of  Water  and  the  Alcohols 213 

Mixtures  of  Acetone  with  Water  and  the  Alcohols 215 

Viscosity  and  Atomic  Volume 217 

Dissociation  in  Nonaqueous  Solvents 219 

Ternary  Mixtures  of  the  Alcohols  with  Water 220 

Glycerol  and  Mixtures  with  the  Alcohols  and  Water 221 

Rubidium  Salts  in  Mixtures  of  Acetone  and  Water 224 

Work  in  Formamide 224 

Viscosities  of  Caesium  Salts 225 

CHAPTER  XII 

COLLOIDAL  SOLUTIONS 

Historical  Sketch 226 

Work  of  Graham 227 

Diffusion  Experiments  of  Graham 227 

Graham's  Views  on  Colloids , 229 

Work  of  M.  Carey  Lea 230 

Nomenclature  Terms  Introduced 231 

Nomenclature  of  Wolfgang  Ostwald 233 

Methods  of  Preparing  Colloidal  Solutions 234 

Double  Decompositions 234 

Reductions 235 

Electrical  Methods 236 

Properties  of  Colloidal  Solutions 237 

Osmotic  Pressures  of  Emulsions  and  of  Suspensions 238 

Lowering  of  Freezing-point  and  of  Vapor-tension 239 

Diffusion  of  Colloids 240 

Brownian  Movement 241 

Distance  Traveled  and  Velocities  of  the  Particles 242 

Cause  of  the  Brownian  Movement 243 

Work  of  Perrin 243 

Concentrations  of  the  Colloids  at  Different  Levels 245 


CONTENTS  xxi 

Radii  of  the  Particles 245 

Mass  of  the  Atom 246 

The  Ultramicroscope 247 

Principle  of  the  Ultramicroscope 247 

The  Zsigmondy  Apparatus 248 

The  Siedentopf  Ultramicroscope  249 

Nomenclature  of  Ultramicroscope 249 

Results  Obtained  with  the  Ultramicroscope.    Sizes  of  the  Colloidal  Par- 
ticles   249 

Motion  of  Emulsion  Particles 251 

Electrical  Properties  of  Colloids 252 

Electrical  Endosmosis 253 

Cataphoresis 254 

Velocities  with  which  the  Particles  Move >. 256 

Cataphoresis  with  Emulsoids 256 

Electrical  Properties  of  Colloids  in  Nonaqueous  Solvents 257 

Electrostenolysis 258 

Conductivity  of  Suspensions 258 

Electrical  Properties  of  Emulsions 259 

Precipitation  of  Colloids  by  Electrolytes 260 

Action  of  Electrolytes  on  Colloidal  Suspensions 261 

Microscopic  and  Ultramicroscopic  Observation  of  the  Precipitation  of 

CoUoids 262 

Cause  of  the  Precipitation 263 

Why  a  Colloid  is  Unstable 263 

How  Electrolytes  Act 264 

Precipitation  and  the  Valency  of  the  Ion  with  Opposite  Charge 265 

Precipitation  of  a  Positive  Sol 265 

Precipitation  of  a  Negative  Sol 266 

Precipitation  not  Proportional  to  Valence 268 

Precipitation  of  Coarser  Suspensions  by  Electrolytes 269 

Precipitation  of  Emulsoids  by  Electrolytes 270 

Action  of  One  Colloid  on  Another 271 

Action  of  Emulsoid  on  Emulsoid 273 

Action  of  Suspensions  on  Emulsions 273 

Protective  Action  of  Colloids 273 

Gels 274 

Physical  Properties  of  Gels 275 

Solventation  and  Desolventation  of  Gels 276 

Less  Elastic  Gels 276 

More  Elastic  Gels 277 

Rate  at  which  Water  is  Taken  up 278 

Heat  Set  Free  in  Imbibition 279 

Imbibition  in  Aqueous  Solutions 280 

Theories  of  CoUoids 280 

Are  Colloids  Solutions? 280 

The  Adsorption  Theory  of  Colloids 282 

The  Suspension  Theory  of_Colloids 282 


xxii  CONTENTS 

Bearing  of  Colloidal  Chemistry  on  other  Branches  of  Science 283 

Bearing  of  Colloids  on  Physiological  Chemistry 284 

Relations  between  Organic  and  Inorganic  Ferments 284 

Action  of  Toxins  on  Antitoxins 286 

Bearing  of  Colloids  on  Pharmacology  and  Pathology 287 

Bearing  of  Colloids  on  Agricultural  Chemistry 288 

Bearing  of  Colloids  on  Mineralogy 289 

Colloidal  Chemistry  and  the  Chemical  Industries 289 

Soaps 290 

Tanning 291 

Dyeing 291 

Other  Technical  Applications  of  Colloids 293 

Adsorption 294 

Nature  of  the  Adsorbent 294 

Effect  of  Temperature  on  Adsorption 296 

Theories  of  Adsorption 297 


CHAPTER  XIII 

SOLUTIONS  IN  SOLIDS  AS  SOLVENTS 

Solutions  of  Solids  Dissolved  in  Solids,  or  Solid  Solutions  Proper 299 

Diffusion  in  Solids 302 

Lowering  of  the  Vapor-tension  of  Solids  by  Other  Solids 303 

Lowering  of  the  Freezing-point  of  a  Solid  by  Another  Solid 304 


CHAPTER  XIV 
THE  NEWER  HYDRATE  THEORY 

Earlier  Work 306 

Relation  between  Lowering  of  the  Freezing-point  of  Water  and  Water  of 

Crystallization  of  the  Dissolved  Substance 307 

Approximate  Composition  of  the  Hydrates  Formed  by  Various  Substances 

in  Solution 310 

Relation  between  Water  of  Crystallization  and  Temperature  of  Crys- 
tallization    314 

Dissociation  as  Measured  by  the  Freezing-point  Method  and  by  the 

Conductivity  Method 315 

Temperature  Coefficients  of  Conductivity  and  Hydration 316 

Coefficients  for  Slightly  Hydrated  Salts 318 

Coefficients  for  Strongly  Hydrated  Salts 319 

Relations  between  the  Coefficients 320 

Relation  between  the  Hydration  of  the  Ions  and  then*  Ionic  Volumes  . . .  322 

Hydration  of  the  Ions  and  the  Velocities  with  which  they  Move 324 


CONTENTS  xxiii 

CHAPTER  XV 

THE  SOLVATE  THEORY  OF  SOLUTIONS 

Hydrate  Theory  for  Aqueous  Solutions  Becomes  the  Solvate  Theory  in 

Solutions  in  General 327 

Spectroscopic  Evidence  Bearing  on  the  Solvate  Theory  of  Solution  — 

Work  of  Uhler 330 

"Solvent  Bands"  —  Work  of  Anderson 332 

"Solvent  Bands"  —  Work  of  Strong 335 

Absorption  Spectra  of  Uranium  Compounds 335 

Absorption  Spectra  of  Neodymium  Chloride  in  Various  Solvents 338 

Absorption  Spectra  of  Neodymium  Nitrate  in  Different  Solvents 339 

Transparency  of  Free  and  of  Combined  Water  —  Work  of  Guy 341 

The  Radiomicrometer 342 

Solutions  more  Transparent  than  pure  Water 342 

Work  of  Shaeffer  and  Paulus 344 

Summary  of  the  Lines  of  Evidence  Obtained  in  this  Laboratory,  Bearing 

on  the  Solvate  Theory  of  Solution 345 

How  the  Present  Solvate  Theory  of  Solution  Differs  from  the  Older  Hy- 
drate Theory 348 

Significance  of  the  Solvate  Theory  of  Solution 350 

The  Solvate  Theory  and  the  Theory  of  Electrolytic  Dissociation 351 

Does  the  Solvate  Theory  Help  to  Explain  any  of  the  Apparent  Excep- 
tions to  the  Theory  of  Electrolytic  Dissociation? 352 

Does  the  Solvate  Theory  Aid  in  Explaining  the  Facts  of  Chemistry  in 

General,  and  of  Physical  Chemistry  in  Particular? 355 

Why  the  Nature  of  Solution  Is  of  such  Vital  Importance  not  only  for 

Chemistry,  but  for  Science  in  General 357 

INDEX..  371 


THE   NATUEE  OF   SOLUTION 

CHAPTER  I 

IMPORTANCE  OF  SOLUTION 

THE  object  of  this  introductory  chapter  is  to  call 
attention  to  the  significance  of  that  condition  of  matter 
which  we  describe  as  dissolved.  We  shall  see  that  it  is 
fundamental  not  only  for  chemistry,  but  for  many  other 
branches  of  natural  science. 

Importance  of  Solution  Early  Recognized.  —  The  bear- 
ing of  solution  on  natural  processes  was  early  recognized. 
It  was  clearly  seen  that  without  solution  there  would  be 
no  chemistry.  This  was  summarized  by  the  alchemists 
in  the  terse  generalization,  "Corpora  non  agunt  nisi  soluta" 
or  in  the  equally  concise,  "Menstrua  non  agunt  nisi 
fluida." 

These  generalizations  are  a  little  too  broad  hi  the  light 
of  what  was  known  about  solutions  at  the  tune  when  they 
were  written.  A  "solution"  in  the  days  of  the  alchemists 
was  primarily  a  solution  of  a  solid  in  a  liquid.  Even  if 
solution  was  not  limited  by  them  to  those  systems  which 
result  when  liquids  are  brought  hi  contact  with  solids, 
this  type  of  solutions  was  given  such  a  prominence,  that 
it  was  generally  in  mind  when  the  term  "solution  "was 
used. 

Types  of  Solutions.  —  The  present  use  of  the  term 
"solution"  is  not  only  much  broader  than  that  adopted 
by  the  alchemists,  but  is  far  broader  than  that  employed 
even  forty  years  ago. 


NATURE  OF  SOLUTION 


We  know  matter  in  three  states  of  aggregation  —  solid, 
liquid,  and  gaseous.  Matter  in  every  one  of  these  three 
states  can  be  dissolved  in  matter  of  the  same  state  of 
aggregation  as  itself  and  in  both  of  the  other  states. 
Thus,  we  have  solutions  of  gases  in  gases,  or  mixtures  of 
gases  which  do  not  act  chemically  upon  one  another. 
The  characteristic  here  is  unlimited  solubility,  the  prop- 
erties of  the  mixture  being  the  sum  of  the  properties  of 
the  constituent  gases. 

Liquids  have  a  vapor-tension  in  the  presence  of  a  gas 
which  is  the  same  as  in  a  vacuum.  There  is  therefore 
limited  solubility  of  the  liquid  in  the  gas. 

Solids  have  a  vapor-tension  in  the  presence  of  a  gas 
which  is  a  function  of  the  temperature. 

Solutions  of  gases,  liquids,  and  solids  in  liquids  are  the 
best  and  longest  known  types  of  solutions. 

Gases  dissolve  in  liquids  to  only  a  limited  extent,  the 
amount,  in  keeping  with  Henry's  law,  increasing  with  the 
pressure  to  which  the  gas  is  subjected. 

Liquids  dissolve  in  liquids,  many  of  them  to  an 
unlimited  extent.  Liquids  which,  at  ordinary  tempera- 
tures, have  only  limited  solubility  in  other  liquids,  often 
become  infinitely  soluble  at  more  elevated  temperatures. 
Solids  dissolve  in  liquids  to  a  limited  extent,  the  amount 
for  any  solid  being  a  function  of  the  temperature. 

Solutions  of  gases  and  liquids  in  solids  are  well  known. 
Carbon  dioxide  dissolves  in  charcoal,  hydrogen  in  many 
metals,  etc.,  and  a  large  number  of  liquids  dissolve  in 
many  solid  substances.  One  of  the  newest  and  most 
interesting  types  of  solutions  is  that  of  solid  in  solid. 
Solid  solutions  came  into  prominence  about  twenty-five 
years  ago,  when  it  was  shown,  as  we  shall  see,  that  mixtures 
of  certain  solids  exhibit  all  of  the  properties  of  solutions 
of  liquids  or  solids  in  liquids.  This  will  be  discussed 
later  at  some  length  under  "solid  solutions." 

We  shall  use  the  term  solution  throughout  this  book 
in  the  broad,  modern  sense  referred  to  above. 


IMPORTANCE  OF  SOLUTION  3 

State  of  Aggregation  a  Function  of  Conditions.  —  The 

relations  between  solids,  liquids,  and  gases,  are  now  pretty 
well  understood.  Formerly,  certain  substances  were  known 
only  in  the  solid  state,  others  only  in  the  liquid  condi- 
tion, while  others  still  were  always  gaseous.  We  now 
know  that  the  state  of  aggregation,  of  elementary  sub- 
stances at  least,  is  primarily  a  function  of  the  temperature 
and  pressure  to  which  they  are  subjected,  and  espe- 
cially of  the  temperature  —  a  gas  to  be  liquefied  must 
not  be  above  the  critical  temperature.  Every  one  of  the 
more  common  elements  is  known  as  a  liquid  and  a  gas, 
and  all  except  helium  as  a  solid.  The  reason  why  helium 
has  not  been  solidified  is  its  very  low  freezing-point, 
below  -268°,  and  the  comparatively  small  quantity  thus 
far  obtainable. 

The  history  of  the  liquefaction  of  the  more  resistant 
gases  is  one  of  the  most  fascinating  chapters  hi  modern 
physical  chemistry.  Of  these,  oxygen  was  the  first  to 
succumb.  This  was  accomplished  by  Cailletet,1  on  the 
one  hand,  and  Pictet,2  on  the  other.  The  former  allowed 
highly  compressed  oxygen  cooled  in  liquid  sulphur  dioxide 
•  to  expand  quickly,  when  some  of  the  oxygen  was  liquefied. 
Its  boiling  point  is  — 184°. 

Hydrogen  was  first  liquefied  by  Wroblewski.3  He 
cooled  highly  compressed  hydrogen  in  liquid  oxygen  which 
was  boiled  under  low  pressure.  The  boiling-point  of  liquid 
hydrogen  is  -252°.  When  liquid  hydrogen  was  evaporated 
under  diminished  pressure  it  froze  at  —255°. 

Compressed  helium,  when  cooled  hi  liquid  hydrogen 
and  allowed  to  expand  suddenly,  was  partly  liquefied  at 
—  268°.  As  already  stated,  helium  has  not  yet  been 
solidified.  For  details  in  connection  with  the  liquefaction 
of  gases  some  work4  must  be  consulted  which  deals 

1  Compt.  Rend.,  85,  1217  (1877). 

2  Ibid.,  85,  1214,  1220  (1877). 
«  Ibid.,  98,  149  (1884). 

4  Hardin:  Liquefaction  of  Gases;  also  the  Author's  Elements  of  Physical 
Chemistry,  4th  edition,  pp.  85  to  90.  (The  Macmillan  Co.) 


4  THE  NATURE  OF  SOLUTION 

especially  with  this  subject.  The  liquefaction  of  gases  is, 
then,  a  completed  chapter  of  science. 

The  melting  of  solids  has  been  facilitated  very  greatly 
by  the  electric  furnace.  By  means  of  it  very  high  tem- 
peratures are  obtainable,  and  under  conditions  which  can 
be  readily  worked  with. 

The  question  as  to  whether  a  compound  is  a  solid, 
liquid,  or  gas,  is  not  simply  a  function  of  temperature  and 
pressure.  Another  factor  comes  into  play,  and  that  is 
the  stability  of  the  compound.  Many  compounds  are 
stable  only  in  the  solid  form.  Others  can  be  liquefied, 
but  cannot  be  volatilized  even  under  low  pressure. 

The  object  of  this  paragraph  is  to  call  attention  to  the 
fact  that  the  former  conceptions  of  solids,  liquids,  and  gases 
have  been  considerably  modified  in  the  light  of  recent 
advances;  and  especially  that  this  is  in  keeping  with  the 
recent  developments  in  connection  with  our  conception 
of  solutions.  Formerly,  we  knew  practically  only  solu- 
tions in  liquids,  as  the  solvents.  Now  we  know  solutions 
in  gases,  in  liquids,  and  in  solids  as  solvents.  There  is 
perhaps  a  certain  analogy  in  the  developments  along  these 
two  different  lines. 

Aqueous  Solutions  the  Most  Important.  —  Solutions 
in  water  as  the  solvent  were  early  recognized  as  the 
most  important.  The  old  Greek  philosophers  regarded 
water  as  "the  beginning  of  all  things."  Said  Thales,  "all 
things  have  their  origin  in  water  and  return  unto  the 
same." 

The  father  of  medicine,  Hippocrates,  laid  great  stress 
upon  the  healing  power  of  water,  and  his  successors 
referred  again  and  again  to  the  same  point.  Van  Helmont, 
hi  the  seventeenth  century,  regarded  water  as  the  sub- 
stance of  all  things.  It  was  the  original  element  into  which 
everything  else  could  be  transformed.  In  the  same 
century  Basil  Valentine  was  equally  insistent  upon  the 
importance  of  water  in  chemistry.  Said  he,  "Water  is 
the  mother  of  all  the  metals." 


IMPORTANCE  OF  SOLUTION  5 

This  brings  us  to  Robert  Boyle,  who  also  lived  in  the 
seventeenth  century.  In  his  book,  "Sceptical  Chymist," 
he  lays  stress  upon  the  earlier  view  that  water  is  the 
origin  of  all  things,  and  adds  that  "it  seems  evident 
that  water  may  be  transmuted  into  all  the  other  ele- 
ments." "Not  only  plants,  but  animals  and  minerals 
may  be  produced  out  of  water."  Boyle  thus  went  farther 
in  his  laudation  of  the  importance  of  water  than  any  of 
his  predecessors. 

The  importance  of  solution,  especially  in  water  as  the 
solvent,  was  by  this  tune  generally  recognized.  Solution 
was  the  fundamental  condition.  Here  all  transformations 
took  place. 

Under  these  conditions  it  was  only  natural  that  chem- 
ists should  look  for  a  universal  solvent  —  something  which 
would  dissolve  everything  else. 

Paracelsus,  hi  the  sixteenth  century,  termed  this  uni- 
versal solvent  and  healing  medium  "alkahest."  Paracel- 
sus did  not  describe  the  physical  and  chemical  properties 
of  alkahest  in  general.  This  remained  for  Van  Helmont. 
It  was  a  liquid  substance  resembling  water,  whence  its 
name  "  Ignisaqua" 

The  existence  of  this  hypothetical,  ideal  solvent  was 
soon  called  hi  question.  In  1675  Lemery  regarded  it  as 
simply  imaginary;  and  Robert  Boyle  in  his  above  men- 
tioned work  called  its  existence  into  question. 

Alkahest  was  a  myth,  but  the  conception  of  a  universal 
solvent  led  to  a  large  amount  of  work  with  solutions,  and 
an  extension  of  knowledge  hi  this  field.  Like  so  many 
other  erroneous  hypotheses  it  thus  led  to  important 
results. 

Present  Use  of  the  Term  Solution.  —  We  recognize 
today,  as  did  the  alchemists,  that  aqueous  solutions  are 
the  most  important.  Water  is  the  most  general  solvent 
known  to  man.  Aqueous  solutions  are  the  most  important 
in  chemistry,  and  are  fundamental,  as  we  shall  see,  for  all 
biological  phenomena. 


6  THE  NATURE  OF  SOLUTION 

We  use  the  term  solution  today,  however,  in  the  broad 
sense  in  which  it  has  already  been  defined.  Solutions  in 
solvents  other  than  water  also  play  a  prominent  r61e 
in  chemistry.  Solutions  hi  the  alcohols  and  other  neutral, 
organic  solvents  are  important  not  only  in  pure  chemistry 
but  in  the  chemical  industries  as  well. 

To  judge  of  the  importance  of  that  condition  of  matter 
known  as  dissolved,  we  must  use  "solution"  in  the 
broader  and  more  modern  sense.  We  must  consider 
solutions  of  gases,  of  liquids,  and  of  solids  in  gases,  in 
liquids,  and  in  solids.  With  this  conception  in  mind  let 
us  see  what  is  the  bearing  of  solution  on  chemistry. 

Solution  and  Chemical  Transformation.  —  We  can 
perhaps  best  understand  the  importance  of  solution  for 
chemistry  by  recalling  the  operations  of  qualitative 
analysis.  Most  of  the  reactions  there  involved  take  place 
in  solution,  and  in  aqueous  solution  at  that.  We  dis- 
solve salts  and  bring  their  solutions  together  in  order 
to  have  them  react.  The  dry,  undissolved  solids  would 
not  react;  their  solutions  react  immediately  on  contact. 

In  some  cases  the  reactions  are  carried  out  at  higher 
temperatures,  and  then  the  salts  are  fused  together.  A 
fused  mixture  of  two  or  more  substances  is  just  as  truly  a 
solution  as  cane  sugar  in  water  at  ordinary  temperatures. 
The  definition  of  solution  as  "a  homogeneous  mixture 
of  two  or  more  substances,  the  constituents  of  which 
cannot  be  separated  mechanically,"  does  not  contain  any 
reference  to  temperature.  We  have  solutions  at  very  low 
temperatures,  at  ordinary  temperatures,  at  elevated  tem- 
peratures, and  at  very  high  temperatures.  They  are  all 
solutions  and  obey  the  laws  of  this  condition  of  matter. 

When  we  want  an  acid  or  a  base  to  react,  we  do  not 
use  the  pure,  homogeneous  substance,  but  we  dissolve 
it  in  water.  As  we  shall  see,  no  pure  acid  or  base,  so-called, 
has  any  acid  or  basic  properties.  They  become  acids 
or  bases  only  when  dissolved  in  water  or  some  other 
solvent.  The  reason  for  this  will  be  pointed  out  later. 


IMPORTANCE  OF  SOLUTION  7 

When  we  wish  to  analyze  any  solid,  one  of  the  first 
steps  consists  hi  finding  a  solvent  for  it  —  hi  getting  it 
into  solution.  Not  until  it  is  dissolved  do  the  general 
methods  of  analysis  apply  to  it.  The  reason  for  this 
will  also  become  apparent  hi  the  proper  place. 

Solution  Fundamental  for  all  Branches  of  Chemistry.  — 
What  has  been  said  in  reference  to  the  importance  of  solu- 
tion for  qualitative  chemistry,  applies  with  equal  force  to 
other  branches  of  chemistry.  Quantitative  chemistry  is 
absolutely  dependent  upon  solution  both  for  gravimetric 
and  for  volumetric  work.  Precipitations  are,  of  course, 
effected  hi  solutions,  and  hi  all  titrations  both  the  matter 
titrated  and  the  substances  used  to  titrate  are  dissolved. 

When  we  turn  to  organic  chemistry  we  find  that  the 
same  condition  obtains.  Reactions  here,  hi  general,  take 
place  hi  solution  hi  the  broader  sense  of  that  term.  Aque- 
ous solutions  are  frequently  used,  but  solutions  hi  sol- 
vents other  than  water  —  non-aqueous  solutions  —  are 
also  employed. 

Certainly  hi  modern  physical  chemistry  solutions  are  fun- 
damentally important.  Indeed,  those  recent  developments 
,in  chemistry  which  are  classed  under  the  head  of  physical 
chemistry  deal  largely  with  the  properties  of  solutions. 

As  we  shall  see,  it  was  the  partial  recognition  of  the 
real  nature  of  solution,  and  of  the  relation  between  solu- 
tions and  gases,  which  made  possible  many  of  these  recent 
developments.  Solution  is  the  keystone  to  chemistry.  We 
might  almost  say  " without  solution,  no  chemistry"  —  at 
least  no  chemistry  hi  the  sense  in  which  we  now  know 
this  large  and  important  branch  of  natural  science. 

Dry  Substances  do  not  Generally  React  Chemically.  — 
The  bearing  of  solution  on  chemistry  can  perhaps  be  seen 
best,  if  we  examine  the  chemical  behavior  of  substances 
when  out  of  solution.  Take  those  substances  whose 
aqueous  solutions  are  chemically  among  the  most  active, 
and  see  how  these  same  substances  behave  when  "dry." 
Here  we  are  using  the  term  "dry,"  not  meaning  simple 


8  THE  NATURE  OF  SOLUTION 

freedom  from  water  or  aqueous  vapor,  but  from  all  dis- 
sociating solvents. 

Dry  Chlorine  and  Dry  Sodium.  —  It  was  shown  by 
Wanklyn1  that  when  dry  chlorine  gas  is  passed  over 
fused  metallic  sodium,  there  is  no  action.  In  his  own 
words,  "When  chlorine  gas  is  passed  over  metallic  sodium 
—  even  when  the  metal  is  fused  and  whilst  in  a  state  of 
fusion,  shaken  in  contact  with  the  gas,  so  as  to  expose 
fresh  metallic  surfaces  —  there  is  no  action.  A  glass 
vessel  containing  a  piece  of  sodium  was  weighed,  and, 
after  the  transmission  of  chlorine  under  the  circumstances 
above  named,  it  was  re-weighed. 

Grams 

Weight  before  the  action  of  chlorine 7.847 

Weight  after  the  action  of  chlorine 7.863 

Gain 0.016 

The  quantity  of  metallic  sodium  taken  for  the  experiment 
was  0.770  grams." 

This  explains  the  result  which  has  so  often  been 
obtained  upon  the  lecture  table,  when  chlorine  gas  was 
carefully  washed  and  dried,  and  then  passed  over  heated 
metallic  sodium  to  show  the  direct  union  of  sodium  and 
chlorine,  forming  sodium  chloride.  The  fused  sodium 
maintained  an  untarnished  surface  in  contact  with  the 
chlorine  gas,  until  the  experimenter,  perhaps  in  disgust, 
opened  the  tube  containing  the  fused  sodium  and  the 
chlorine  gas.  On  opening  the  tube  a  little  moist  air 
was  introduced;  enough  moisture  being  added  to  cause 
the  sodium  and  chlorine  to  react  with  a  violence  which 
gave  more  satisfaction  to  the  observing  class  of  beginners 
than  to  the  experimenter. 

Dry  Hydrochloric  Acid  Gas  on  Dry  Carbonates. — 
The  action  of  dry,  hydrochloric  acid  gas  on  Iceland  spar 
and  witherite  was  studied  by  Hughes.2  In  his  own  words, 
"  The  variations  in  weight  which  we  observed  were  so 

*  Chem.  News.  20,  271  (1869). 
8  Phil  Mag.,  34,  117  (1892). 


IMPORTANCE  OF  SOLUTION  9 

minute  that  no  definite  assertion  can  be  made  of  an  action 
having  taken  place;  and  the  slight  variations  observed  may 
be  due  to  experimental  errors,  and  to  the  imprisonment  or 
entanglement  of  the  molecules  of  hydrochloric  acid  gas 
amongst  the  finely  divided  particles  of  the  Iceland  spar  or 
witherite." 

Dry  Acids  on  Dry  Litmus.  —  A  still  more  remarkable 
experiment  was  described  by  Marsh.1  "An  acid  to  affect 
blue  litmus  must  be  more  or  less  dilute;  hi  other  words, 
water  is  necessary  for  the  particular  action,  whatever 
it  may  be,  which  occurs  when  blue  litmus  becomes  red." 
"  I  have  found  that  dry  litmus  paper  in  hydrochloric 
acid  dried  with  phosphorus  pentoxide,  does  not  appre- 
ciably alter  in  color  for  some  time.  Ordinary  concen- 
trated sulphuric  acid  does  not  redden  litmus  paper,  but 
imparts  a  more  or  less  bluish  purple  tint  to  it;  and  the 
Nordhausen  acid  acts  similarly." 

Dry  Ammonium  Chloride  Sublimes  Undecomposed. — 
More  remarkable,  however,  than  any  of  the  above  experi- 
ments is  that  carried  out  by  Baker.2  His  paper,  "In- 
fluence of  Moisture  on  Chemical  Change,"  is  one  of  the 
most  interesting  that  has  ever  been  written  upon  this 
subject.  He  points  out  that  "dried  ammonium  chloride 
may  be  sublimed  from  a  mixture  of  the  salt  with  dried  lime 
without  ammonia  being  liberated";  and  that  "dried 
ammonium  chloride  does  not  dissociate  when  heated  to 
350°." 

Dry  Hydrochloric  Acid  Gas  on  Dry  Ammonia.  —  The 
experiment  described  in  this  paper  by  Baker,  upon  which  it 
is  desired  to  lay  special  stress,  is  the  behavior  of  a  mixture 
of  dry  hydrochloric  acid  gas  and  dry  ammonia.  Helm- 
holtz  and  Richartz  had  noted  that  no  fumes  are  pro- 
duced when  these  gases,  previously  dried,  are  brought 
together;  they,  however,  believed  that  the  chemical  action 
still  went  on,  although  the  ammonium  chloride  was  not 
precipitated  in  the  form  of  powder. 

1  Chem.  News,  61,  2  (1890).  »  Journ.  Chem.  Soc.,  65,  611  (1894). 


10 


THE  NATURE  OF  SOLUTION 


Baker  prepared  his  ammonia  by  heating  purified 
ammonium  chloride  with  purified  lime.  He  passed  the 
ammonia  first  over  solid  potash,  and  then  over  a  mixture 
of  copper  oxide  and  potassium  oxide.  The  latter  was 
recommended  by  Stas  as  being  a  better  drying  agent  than 
phosphorus  pentoxide.  Baker  prepared  it  by  fusing 
finely  divided  copper  with  potassium  nitrate.  The  am- 
monia was  then  introduced  into  one  arm  of  the  tube 
[Fig.  1],  over  phosphorus  pentoxide.  Ammonia  thus  pre- 
pared had  no  action  on  phosphorus  pentoxide. 


HCl 


J  Mercury 
FIG.  1. 

The  hydrochloric  acid  gas  w£,s  prepared  by  the  action 
of  pure  sulphuric  acid  on  recrystallized  sodium  chloride. 
The  gas  was  introduced  into  the  other  arm  of  the  tube,  and 
allowed  to  stand  for  a  week  in  contact  with  phosphorus 
pentoxide. 

When  the  gases  were  mixed  no  white  fumes  were  pro- 
duced. "  On  connecting  the  interior  of  the  tube  with  dried 
mercury,  by  means  of  the  third  limb  of  the  tap,  the  mer- 
cury did  not  rise,  showing  that  no  combination  of  the 
gases  had  taken  place."  "If  a  trace  of  moist  air  be  ad- 
mitted to  the  mixture  of  dry  gases,  dense  white  fumes 
are  at  once  produced,  and  the  mercury  rushes  up  into  the 
tube." 

Gutmann1  obtained  opposite  results  from  Baker,  but 
Baker2  repeated  his  work  and  showed  that  Gutmann  had 
not  taken  the  proper  precautions  in  drying  the  gases. 
Gutmann  used  phosphorus  pentoxide  which  probably  con- 


IMPORTANCE  OF  SOLUTION  11 

tained  metaphosphoric  acid;  if  so,  this  was  the  cause  of  the 
different  result  obtained  by  him.  In  his  later  work,  Baker 
obtained  the  same  result  with  reference  to  the  inaction  of 
dry  ammonia  and  dry  hydrochloric  acid  gas,  and  also 
found  that  properly  dried  ammonium  chloride  when  vola- 
tilized had  a  normal  molecular  weight. 

Dry  Sulphuric  Acid  on  Dry  Metallic  Sodium.  —  A 
result  still  more  remarkable  than  any  of  those  already 
referred  to,  is  the  inactivity  of  dry  sulphuric  acid  on 
dry  metallic  sodium.  If  there  is  any  reaction  in  all 
chemistry  which  we  would  suppose,  on  a  priori  ground, 
must  always  take  place  whenever  the  constituents  were 
brought  together,  it  would  be  the  reaction  between  metal- 
lic sodium  and  sulphuric  acid.  From  our  general  knowl- 
edge of  the  substances  chemically,  it  is  difficult,  not  to 
say  impossible,  to  conceive  of  them  existing  hi  the  pres- 
ence of  one  another  without  entering  into  chemical 
combination. 

The  writer  has  seen  the  folio  whig  experiment3  per- 
formed. A  piece  of  metallic  sodium  which  had  been  dried 
with  special  care  over  phosphorus  pentoxide,  was  immersed 
in  sulphuric  acid  which  had  also  been  dried  with  very 
special  care.  .  When  the  sodium  first  touched  the  sul- 
phuric acid  there  was  incipient  reaction,  due  to  the  sodium 
having  taken  up  on  its  surface  a  little  moisture  hi  passing 
through  the  air  during  the  transfer  from  one  vessel  to 
another.  This  was  over  hi  a  moment.  The  sodium  then 
remained  suspended  in  the  sulphuric  acid  as  quietly  as  if 
hi  ligroin. 

If  one  desires  to  repeat  any  of  the  above  described 
experiments,  very  special  precautions  must  be  taken  to  dry 
all  of  the  substances  involved.  The  difficulty  in  removing 
water  from  almost  anything  can  be  appreciated  when  we 
consider  that  Stas  was  able  to  pump  water  out  of  a  glass 
tube  which  had  been  heated  to  redness  for  a  month,  the 

1  Lieb.  Ann.,  299,  267  (1898).  2  Journ.  Chem.  Soc.,  73,  422  (1898). 

3  Proceed.  Chem.  Soc.,  p.  86  (1894). 


12  THE  NATURE  OF  SOLUTION 

tube  having  been  attached  all  the  while  to  a  vacuum  pump. 
This  shows  that  the  usual  methods  of  drying  are  entirely 
inadequate  in  work  such  as  that  which  has  just  been  de- 
scribed. The  failure  to  recognize  this  fact  has  led  certain 
experimenters  to  perform  some  of  the  above  experiments  to 
their  sorrow. 

What  has  been  stated  in  reference  to  the  bearing  of 
solution  on  chemical  action  could  be  elaborated  and  sup- 
plemented almost  indefinitely.  Enough  has  been  given, 
however,  to  show  that  things  in  the  pure,  homogeneous 
condition  do  not,  in  general,  react  chemically.  It  is  only 
when  matter  in  one  state  of  aggregation  is  mixed  with 
matter  in  the  same  or  in  a  different  state  of  aggregation 
—  is  dissolved  —  that  it  becomes  really  active  from  the 
chemical  standpoint.  The  melting  of  a  solid  or  the  vola- 
tilization of  a  liquid  are  not  ordinarily  considered  as 
chemical  reactions. 

Dry  substances  not  infrequently  react  when  fused  to- 
gether at  elevated  temperatures.  In  such  solutions  rise 
in  temperature  produces  an  effect,  as  we  shall  see,  anal- 
ogous to  that  of  the  solvent  at  ordinary  temperatures. 
This  is  in  no  wise  opposed  to  the  point  which  is  brought 
out  in  the  preceding  sections.  Of  chemical  reactions  in  gen- 
eral, relatively  few  take  place  out  of  solution.  Solution 
is  therefore  absolutely  fundamental  for  chemistry. 

Importance  of  Solution  for  Physics.  —  When  we  turn 
from  chemistry  to  physics  we  find  solution  not  playing  so 
prominent  a  role  as  in  chemistry,  but  still  one  of  impor- 
tance. The  conduction  of  heat  in  solutions,  the  absorption 
of  light  by  solutions,  and  the  passage  of  electricity  through 
solutions  are  important  for  physics.  We  may  also  mention 
surface-tension,  viscosity,  diffusion,  polarization,  electrolysis 
etc.,  of  solutions;  phenomena  which  have  long  attracted 
the  attention  of  physicists.  Again,  consider  the  primary 
and  secondary  electrical  batteries.  The  primary  cells, 
prior  to  the  invention  of  the  dynamo,  were  for  a  long  time 
the  chief  source  of  electricity;  and  primary  cells  in  general 


IMPORTANCE  OF  SOLUTION  13 

depend  for  their  existence  upon  the  solutions  of  the  elec- 
trolytes around  their  poles.  Secondary  cells  are  equally 
dependent  upon  solutions,  and  secondary  cells  are  coining 
more  and  more  into  prominence  hi  studying  electrical 
phenomena.  These  are  only  a  few  of  the  many  applica- 
tions of  solutions  in  physics. 

Solution  and  Geology.  —  Solution,  used  hi  the  broad 
sense  already  indicated,  is  of  importance  hi  geology. 
Rocks  in  general  are  formed  from  molten  magmas  (igneous 
rocks),  or  are  deposited  from  aqueous  solutions  or  sus- 
pensions in  water  (sedimentary  rocks);  the  molten  mag- 
mas being  simply  solutions  of  various  substances  hi  one 
another  at  high  temperatures;  and  whether  a  system  is, 
or  is  not,  a  solution  does  not  raise  the  question  of  tem- 
perature. As  we  have  seen,  we  can  have  solution  at  any 
temperature.  A  molten  geological  magma  at  a  tempera- 
ture of  several  thousand  degrees,  is  as  much  a  solution 
as  alcohol  or  cane  sugar  hi  water  at  ordinary  tempera- 
ture. Solution  is,  therefore,  fundamental  to  this  branch 
of  geology. 

The  sedimentary  rocks  are  deposited  from  aqueous 
suspensions  or  aqueous  solutions.  The  solid  material  is 
often  in  a  state  of  purely  mechanical  suspension  hi  the 
water.  In  other  cases  it  is  more  finely  divided,  or  hi  a 
state  of  colloidal  suspension,  as  it  is  termed.  It  may  be 
still  more  finely  divided  and  exist  as  a  colloidal  solution; 
or  it  may  be  in  the  water  as  a  true  solution.  In  any  case 
the  liquid  is  essential  to  the  formation  of  the  sedimentary 
rocks.  This  does  not  refer  to  the  aeolian  or  to  the  glacial 
rocks. 

Minerals  crystallize  in  general  either  from  molten 
masses  or  from  aqueous  solutions  —  in  either  case  from 
solution  —  and  mineralogy  is  an  important  branch  of 
geology. 

Heinrich  Rose  called  attention  to  a  geological  action 
which  is  now  generally  recognized,  illustrating  the  impor- 
tant role  of  solution  in  geology.  Carbon  dioxide  in  the  air 


14  THE  NATURE  OF  SOLUTION 

or  in  the  water  in  the  earth,  acting  through  long  periods 
of  time,  decomposes  the  silicates.  This  is  an  important 
factor  in  the  "weathering"  of  the  rocks.  Geologically, 
weathering  is  an  important  phenomenon.  Over  the  sur- 
face of  the  globe  we  have  the  stable  silicates  being  con- 
verted into  the  less  stable  carbonates  —  the  carbonation  of 
the  rocks,  and  this  materially  alters  the  face  of  the  globe. 

This  action  of  large  amounts  of  carbon  dioxide  over 
long  periods  of  time  illustrates  a  principle  which  is  also 
of  fundamental  significance  for  chemistry.  The  silicates 
are  among  the  most  stable  and  insoluble  compounds. 
Carbonic  acid  —  carbon  dioxide  in  water  —  is  one  of  the 
weakest  acids.  Yet  this  very  weak  acid,  acting  in  enor- 
mous quantities,  which  is  what  the  above  amounts  to,  is 
capable  of  decomposing  large  amounts  of  the  silicates. 
This  is  one  of  the  best  examples,  at  least  in  nature,  of  the 
effect  of  mass  or  quantity  as  conditioning  the  magnitude 
and  even  the  direction  of  a  chemical  reaction.  After  Rose 
called  attention  to  the  wide-reaching  significance  of  this 
reaction  not  only  for  geology,  but  for  chemistry,  no  one 
doubted  the  importance  of  the  role  played  by  mass  in 
chemistry.  Indeed,  this  was  one  of  the  most  important 
observations  made  bearing  upon  the  effect  of  mass,  and  it 
played  an  important  part  in  calling  the  attention  of 
chemists  to  the  fact  that  they  must  take  into  account  not 
only  the  nature  of  the  reacting  substances,  but  the  relative 
quantities  which  were  brought  together.  It  is  an  impor- 
tant chapter  in  the  history  of  the  development  of  the  law 
of  mass  action.1 

Solution  and  Biology.  —  Solution  is  quite  as  important 
for  the  various  branches  of  biology  as  it  is  for  chemistry. 
Take  zoology.  Animals  are  composed  largely  of  water.  If 
we  dessicate  almost  any  animal,  it  loses  more  than  half 
its  weight,  and  most  animals  lose  much  more  than  half 
their  weight.  Take  a  human  being  of  average  weight, 

1  See  the  Author's  Elements  of  Physical  Chemistry.  4th  edition.  (The  Mac- 
millan  Co.) 


IMPORTANCE  OF  SOLUTION  15 

say  about  one  hundred  and  fifty  pounds;  when  com- 
pletely desiccated  he  would  weigh  about  sixty  pounds. 

More  important  than  the  large  preponderance  of  water 
in  animals  is  its  function.  Animals,  hi  general,  cannot 
"live"  long  without  water,  using  the  term  "live"  in  the 
sense  of  grow  and  reproduce.  An  amoeba  may  be  roughly 
described  as  a  solution  of  a  large  number  of  substances 
in  water,  surrounded  by  a  membrane,  and  what  is  said  of 
amoeba  applies  to  cells  in  general,  and  tissues  of  animals 
are  made  up  of  cells.  Remove  the  solvent,  and  the  animal 
either  dies  or  lies  dormant,  and  in  the  latter  case  may 
exhibit  the  phenomena  of  "life"  again  when  the  removed 
water  is  restored  to  it. 

Again,  animals  obtain  most  of  then*  food  either  hi  solu- 
tion outside  of  their  body,  or  the  food  passes  into  solu- 
tion after  it  is  taken  into  the  body.  In  either  case  the 
food  is  in  solution,  true  or  colloidal,  or  hi  mechanical  sus- 
pension in  water  before  it  is  assimilated  by  the  animal. 

Solution  is  no  less  important  in  physiological  botany 
than  in  zoology.  Vegetable  life  depends  upon  water  for 
its  existence.  This  is  readily  seen  in  a  gross  way  by 
comparing  the  vegetable  life  in  sections  which  are  well 
watered,  with  that  of  the  desert,  and  even  hi  most  des- 
erts there  is  some  water.  Plants,  like  animals,  are  made  up 
of  cells,  and  plant-cells,  like  the  cells  of  annuals,  are  mainly 
aqueous  solutions. 

Plants  as  well  as  annuals  obtain  their  food  largely  hi 
solution  hi  water.  This  is  dissolved  from  the  soil  and  air 
and  then  assimilated  by  the  plant.  Solution  is  thus 
fundamental  to  botany. 

Among  the  newest  of  the  biological  sciences  is  bac- 
teriology. Without  raising  the  question,  which  is  purely 
one  of  definition,  as  to  whether  these  microscopic  organ- 
isms belong  hi  the  classification  animal  or  plant,  let  us 
see  what  they  are. 

These  relatively  simple  forms  of  life  are  also  composed 
of  cells.  A  bacterium,  from  our  standpoint,  may  be  re- 


16  THE  NATURE  OF  SOLUTION 

garded  as  a  very  complex  aqueous  solution,  surrounded  by 
a  membrane.  These  simple  forms  of  life,  like  the  more 
complex,  depend  for  their  existence  upon  solution.  Indeed, 
they  may  be  said  to  be  primarily  solutions  existing  under 
certain  very  special  conditions. 

We  now  turn  to  the  science  of  pharmacology.  Here 
again  solutions  are  fundamental.  Drugs  are  taken  into  the 
body  either  in  solution,  or  pass  into  solution  in  the  juices 
of  the  body,  which  are  essentially  aqueous  solutions,  be- 
fore they  are  functional;  and  finally  consider  physiological 
chemistry.  The  reactions  here  are  almost  as  much  a 
matter  of  solution,  true  or  colloidal,  as  in  any  other 
branch  of  chemical  science. 

All  in  all,  we  can  see  from  this  very  brief  sketch  that 
water  is  absolutely  essential  to  life.  Its  chief  r61e  in 
living  processes  is  that  of  solvent.  Aqueous  solutions  are, 
then,  absolutely  essential  to  life.  We  often  say  "without 
carbon,  no  life."  We  could  just  as  truly  state  "without 
water,  no  life,"  at  least  in  its  normal  condition. 

Water  a  Remarkable  Compound.  —  We  are  so  familiar 
with  water  that  we  are  liable  to  look  upon  it  as  a  com- 
pound of  relatively  little  significance.  It  is  formed  in 
such  a  large  percentage  of  chemical  reactions,  and  we 
express  it  in  so  many  of  our  chemical  equations,  that  we 
are  in  danger  of  losing  sight  of,  or  never  thinking  of,  its 
importance. 

Having  seen  a  little  of  the  fundamental  significance 
of  aqueous  solutions  for  chemistry,  physics,  geology,  and 
biology,  perhaps  it  would  not  be  too  great  a  digression 
to  consider  very  briefly  a  few  of  the  more  striking 
peculiarities  presented  by  water  itself. 

In  the  first  place  its  composition  is  remarkable.  It  is 
composed  of  the  two  most  important  elements  —  oxygen 
and  hydrogen.  These  are  the  two  elements  which  are  re- 
quired for  the  formation  of  acids  and  bases.  The  con- 
stituent of  all  acids  —  that  which  gives  us  acidity,  is 
hydrogen,  under  a  certain  condition,  as  we  shall  see,  i.e., 


IMPORTANCE  OF  SOLUTION  17 

carrying  a  positive  charge  of  electricity.  The  character- 
istic constituent  of  bases  is  one  hydrogen  atom  united 
with  one  oxygen  atom,  forming  what  is  known  as  the 
hydroxyl  group  and  this  group  carrying  one  negative 
charge  of  electricity,  as  we  shall  also  see.  Water  is 
formed  by  the  union  of  the  hydroxyl  of  the  base  with 
the  hydrogen  of  the  acid  and  acids  and  bases  are  of  the 
utmost  importance  in  our  study  of  chemistry. 

In  addition  to  the  neutralization  of  acids  by  bases, 
think  of  how  many  reactions  there  are  hi  chemistry  hi 
which  water  is  formed.  One  of  the  substances  contains 
hydrogen  and  the  other  hydroxyl,  and  when  they  are 
brought  together  water  is  formed  and  the  reaction  takes 
place.  This  is  true  especially  of  compounds  of  carbon. 
It  is  safe  to  say  that  a  very  large  percentage  of  all  the 
chemical  reactions  known,  take  place  on  account  of  the 
formation  of  water  when  the  reacting  substances  are 
brought  together,  and  it  will  be  shown  that  whenever 
electrically  charged  hydrogen  and  hydroxyl  are  brought 
together  in  appreciable  quantities,  they  combine.  This 
shows  the  importance  of  the  water  in  effecting  so  many  of 
our  chemical  reactions. 

Heat  Evolved  in  the  Formation  of  Water.  —  When 
water  is  formed  by  the  union  of  hydrogen  with  oxygen, 
we  have  the  most  exothermic  or  heat-evolving  chemical 
reaction  known.  When  hydrogen  and  oxygen  combine 
they  liberate  more  heat,  per  equivalent  quantities,  than 
any  other  reaction  in  the  whole  field  of  chemistry.  The 
resulting  compound  —  water  —  is  one  of  the  most  stable 
of  chemical  compounds,  just  as  we  would  expect  it  to  be. 
Those  compounds  formed  with  the  largest  heat  evolution, 
other  things  being  equal,  are  the  most  stable. 

Physical  Properties.  —  Given  the  compound,  water, 
its  physical  properties  are  in  general  extreme  —  their  nu- 
merical expressions  are  either  extremely  large  or  extremely 
small,  and  usually  extremely  large.  Its  specific  heat  is 
among  the  largest  for  any  liquid  substance.  It  is  the 


18  THE  NATURE  OF  SOLUTION 

best  solvent  of  all  known  liquids.  It  has  the  highest 
dielectric  constant  or  specific  inductive  capacity  of  any  of 
the  more  common  liquids.  This  means  that  two  elec- 
trical charges  separated  by  water  have  a  smaller  attraction 
for  one  another  than  when  separated  by  any  of  the  other 
common  liquids. 

As  we  shall  see,  the  power  of  liquid  solvents  to  break 
molecules  of  acids,  bases,  and  salts  down  into  charged 
parts  or  ions  as  they  are  termed,  is  nearly  proportional 
to  the  dielectric  constants  of  the  liquids.  Water,  of  all 
the  common  liquid  solvents,  should  therefore  be  the  best 
dissociant,  and  such  is  the  fact. 

One  other  property  of  water,  on  account  of  its  impor- 
tance in  the  economy  of  nature,  calls  for  special  comment. 
When  most  liquids  are  cooled  they  contract  until  their 
freezing-point  is  reached,  when  they  begin  to  solidify. 
Not  so  with  water.  It  contracts  to  a  temperature  of  4°  C. 
and  then,  with  further  cooling,  expands  until  its  freezing- 
point  is  reached.  A  few  other  liquids  are  known  which 
behave  in  a  similar  manner.  This  may  seem  to  be  an 
accidental  property  of  water  and  of  no  very  great  sig- 
nificance; but  a  moment's  thought  will  show  that  such 
is  not  the  case.  Did  water  contract  until  it  froze,  the 
order  and  economy  of  nature  would  in  a  short  tune  be 
greatly  changed.  The  fact  that  water  expands  from  4°C. 
to  0°  is  the  cause  of  water  freezing  first  on  the  surface  and 
not  at  the  bottom.  Did  water  freeze  first  on  the  bottom 
of  our  rivers,  lakes,  and  larger  bodies  of  water,  fresh  water 
life  in  our  northern  climates  would  soon  be  exterminated. 
Those  forms  of  life  which  depend  for  food  upon  these 
fresh  water  forms,  would  be  compelled  to  migrate  or 
would  die,  and  thus  the  present  distribution  or  equilibrium 
between  the  various  living  forms  would  be  destroyed. 
This  would,  of  course,  seriously  disturb  the  existing  order 
of  things. 

These  are  a  few,  and  a  very  few  of  the  unique  properties 
of  water.  To  discuss  them  at  all  fully  would  require  a 


IMPORTANCE  OF  SOLUTION  19 

separate  volume.  They  suffice,  however,  to  illustrate  the 
point  that  water  is  not  only  a  remarkable  substance,  but, 
all  things  considered,  it  is  by  far  the  most  remarkable  and 
important  chemical  compound  known  to  man. 

Importance  of  Solution  Justifies  the  Work  Expended 
upon  It.  —  There  are  few,  if  any,  subjects  hi  chemistry 
upon  which  so  much  work  and  thought  have  been  ex- 
pended, especially  hi  recent  times,  as  upon  the  nature 
of  solution.  The  reason  for  this  ought  now  to  be  apparent. 
In  order  that  matter  should  react  chemically  it  must  be 
dissolved.  Matter  hi  the  pure  homogeneous  condition, 
as  we  have  seen,  is  comparatively  inert.  Chemistry 
depends  for  its  existence  upon  solution.  The  physical, 
geological,  and  biological  action  of  solutions  is  closely  con- 
nected with  the  chemistry  of  solution. 

In  solution  we  are  dealing,  then,  not  only  with  a  fun- 
damental condition  of  matter,  but  with  the  fundamental 
condition.  This  is  why  so  much  stress  has  been  laid  upon 
solution  in  the  past,  and  interest  in  its  nature  will  prob- 
ably continue  as  long  as  the  natural  sciences  are  studied. 


CHAPTER  II 

EARLIER  VIEWS  AS  TO  THE  NATURE  OF  SOLUTION 

THE  views  of  the  Greek  philosophers  as  to  the  nature 
of  matter  in  general,  and  solution  in  particular,  can  be 
learned  from  a  history  of  chemistry  dealing  with  that 
period.  Much  of  it  is  very  indefinite  and  therefore  does 
not  vitally  concern  us  today. 

A  good  part  of  the  speculation  of  the  earlier  alchemists 
had  to  do  with  the  modus  operandi  of  solution,  rather 
than  with  the  nature  of  solution  itself.  There  were  those 
who  held  that  the  dissolved  particles  simply  went  into  the 
interstices  between  the  molecules  of  the  solvent,  and 
those  who  thought  that  there  was  some  kind  of  a  union 
between  the  solvent  and  the  dissolved  substance. 

Isaac  Newton.  —  It  was  Newton  who,  in  the  seven- 
teenth century,  discovered  the  law  of  gravitation,  that 
matter  attracts  matter  inversely  as  the  square  of  the  dis- 
tance, and  directly  proportional  to  the  product  of  the  masses. 
This  was  found  to  apply  to  matter  at  great  distances  apart, 
and  it  was  only  natural  that  he  should  attempt  to  apply  it 
or  some  other  force  to  matter  removed  from  other  matter 
only  an  infinitesimal  distance,  as  is  the  case  with  the 
atoms.  We  had  what  we  might  term  the  gravitational 
theory  of  chemical  action. 

Newton  extended  this  conception  also  to  solution.  He 
observed  that  when  a  solid  is  dissolved  in  a  liquid  which 
is  much  lighter  than  the  solid,  the  heavier  solid  does  not 
settle  to  the  bottom  of  the  containing  vessel,  but  dis- 
tributes itself  throughout  the  entire  volume  of  the  solvent. 
He  asked  "does  not  this  indicate  that  the  dissolved  parts 
endeavor  to  expand  themselves  and  get  as  far  asunder 
as  the  quantity  of  water  in  which  they  float  will  allow? 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION         21 

And  does  not  this  endeavor  imply  that  they  have  a 
repulsive  force  by  which  they  fly  from  one  another,  or  at 
least,  that  they  attract  the  water  more  strongly  than  they 
do  one  another."1 

This  discovery  of  Newton  marked  a  new  epoch  in 
dealing  with  solutions.  The  solute  dissolved  in  the  solvent 
because  of  an  attraction  between  the  two. 

Boerhaave.  —  The  view  of  Newton  was  somewhat  modi- 
fied and  very  greatly  elaborated  by  Boerhaave.2  "It  fol- 
lows from  the  nature  of  a  solvent,  that  if  it  acts  on  a 
substance  dissolving  it,  hi  a  similar  manner  the  substance 
dissolves  the  solvent.  The  particles  of  the  solvent  and 
those  of  the  dissolved  substance  unite,  after  solution,  to 
form  a  new,  homogeneous  substance."  .  .  .  "The  cause 
of  this  must  be  sought  for  in  both  the  solvent  and  the  dis- 
solved substance.  It  is  common  to  them  both  and  acts 
reciprocally  in  both." 

All  that  was  then  known  of  the  action  of  a  solvent,  was 
that  its  parts  unite  with  those  of  the  dissolved  sub- 
stance. "  We  are  not  dealing  here  with  mechanical 
action,  or  with  violent  repulsion  or  enmity,  but  rather 
with  friendliness,  if  we  can  apply  this  term  to  a  tendency 
to  union."  "  Particulae  solventes  et  solutae  .  .  .  se  affini- 
tate  suae  naturae  colligant  in  corpora  homogenea." 

"The  changes  which  the  solvent  brings  about  hi  the 
dissolved  substance  appear  to  be  due  to  the  intimate 
union  between  the  smallest  particles  of  the  solvent  and  of 
the  dissolved  substance." 

Thus  Newton  and  Boerhaave  placed  the  whole  subject 
of  solution  upon  a  new  basis.  Instead  of  speculating 
about  the  shapes  and  disposition  of  the  dissolved  particles, 
they  attempted  to  answer  the  far  more  fundamental 
questions,  what  is  the  cause  of  solution  and  what  is  its 
nature? 

1  See  Walden:    Die  Losungstheorien  in  ihrer  geschichtlichen  Aufeinander- 
folge,  p.  33. 

2  Ibid.,  p.  33,  from  H.  Boerhaave:  Elementa  Chemia. 


22  THE  NATURE  OF  SOLUTION 

Their  answer  pointed  out  the  analogy  between  solu- 
tion and  chemical  action,  in  that  both  consisted  in  chemical 
union.  They  did  more  than  this,  hi  calling  attention 
to  the  wide-reaching  significance  of  solution  for  chemistry. 
The  result  was  that  a  little  later  we  find  Morveau  asserting 
that  "chemistry  is  primarily  the  science  of  solutions.'' 

Wallerius.  —  The  next  advance  in  the  study  of  solution 
we  owe  to  Wallerius,1  who  raised  the  question,  What  is  the 
cause,  or  what  are  the  causes  of  solution?  Among  these  he 
enumerated  the  following: 

There  must  be  a  similarity  between  the  solvent  and 
solute.  Due  to  this  similarity  there  is  an  attractive 
force  between  the  parts,  which  is  greater  the  more  closely 
the  parts  resemble  one  another.  Due  to  this  attractive 
force  the  solvent  and  solute  unite  with  one  another. 

If  we  omit  the  first  named  cause,  i.e.,  the  similarity 
between  solvent  and  solute,  we  find  these  suggestions  of 
Wallerius  very  important.  While  he  was  not  the  first  to 
point  out  the  existence  of  an  attraction  or  an  affinity 
between  solvent  and  solute,  he  laid  much  stress  upon  it 
and  pointed  out  the  result  of  its  action  —  combination 
between  the  two.  He  also  called  attention  to  the  fact 
that  chemical  combination  is  due  to  an  attraction  or 
affinity  between  the  reacting  substances,  which  is,  of 
course,  a  matter  of  fundamental  importance. 

This  brings  us,  in  the  study  of  solution,  down  to  the 
end  of  the  eighteenth  and  the  beginning  of  the  nine- 
teenth century. 

Lavoisier. —  The  distinguished  French  chemist,  Lavoisier, 
in  1789,  in  his  "Traite*  Elementaire  de  Chimie,"2  made  an 
important  contribution  at  least  to  the  classification  of 
solutions.  "In  solutions  of  salts  the  salt  molecules  are 
simply  torn  apart  from  one  another,  neither  the  salt 
nor  the  water  suffering  any  decomposition.  Both  can 

1  Physische  Chemie. 

2  See  Walden:    Die  Losungstheorien  in  ihrer  geschichtlichen  Aufeinander- 
folge,  p.  47. 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        23 

be  recovered  in  the  same  quantity  as  before  the  opera- 
tion. The  same  can  be  said  of  the  solution  of  the  resins 
in  alcohol  and  in  the  spirituous  solvents.  On  the  other 
hand,  hi  the  dissolution  of  the  metals  there  is  always  a 
decomposition  either  of  the  acid  or  of  the  water.  The 
metal  becomes  oxidized,  passing  into  the  state  of  oxide; 
a  gaseous  substance  is  given  off  —  none  of  the  substances 
after  dissolution  are  in  the  same  state  as  before." 

Lavoisier  thus  distinguished  clearly  between  what  he 
called  solution  and  dissolution;  the  former  being  more 
nearly  what  we  usually  understand  by  the  term  solution, 
the  latter  involving  a  chemical  act.  A  solution,  according 
to  Lavoisier,  is  a  mechanical  mixture. 

Fourcroy.  —  When  we  enter  the  nineteenth  century, 
we  find  Fourcroy1  also  regarding  solution  as  physical,  the 
dissolved  substance  being  simply  hi  a  state  of  fine  mechani- 
cal division  hi  the  solvent. 

Klaproth.  —  Klaproth2  considered  solution  as  the  result 
of  the  action  of  chemical  affinity  between  the  solvent  and 
the  dissolved  substance.  This  means  that  there  is  a 
reciprocal  attraction  between  the  particles  of  the  solute  and 
those  of  the  solvent;  and  this  attraction  must  be  stronger 
than  that  of  the  solute  particles  for  one  another,  or  the 
solvent  particles  for  one  another.  As  he  says,  this  attrac- 
tion between  the  particles  of  the  solute  and  solvent  must 
be  stronger  than  then*  "  cohesion." 

We  now  come  to  the  contribution  to  our  knowledge 
of  solution  by  one  of  the  greatest  French  chemists  of  the 
beginning  of  the  nineteenth  century  —  Berthollet. 

Berthollet.  —  His  views  were  expressed  in  his  great  book, 
"  Essai  de  Statique  Chimique,"  —  the  same  work  hi 
which  the  effect  of  mass  on  chemistry  was  first  so  clearly 
pointed  out.  He  also  called  attention  to  the  fact  that 
solution  is  due  to  a  force  which  must  be  great  enough  to 
overcome  the  cohesion  of  the  dissolved  substance. 

1  Sys&me  des  Connaissances  Chimiques. 

8  Systematisches  Handbuch  der  gesamten  Chemie. 


24  THE  NATURE  OF  SOLUTION 

Stable  compounds  were  called  by  Berthollet,  "  Com- 
binations," the  less  stable  "  Dissolutions."  Solution  is  a 
true  chemical  combination;  the  difference  between  a 
solution  and  a  chemical  compound  is  to  be  found  in  the 
firmness  of  the  union  of  the  parts.  In  a  solution  the  parts 
are  less  firmly  united  and  the  characteristic  properties  of 
dissolved  substances  have  not  been  lost.  Chemical  union 
and  solution  must  therefore  follow  the  same  laws. 

Berthollet,  as  is  well  known,  did  not  believe  in  the  laws 
of  definite  and  multiple  proportions  in  chemistry.  To 
quote  his  own  words  from  this  same  book,1  "Chemists 
maintain  that  in  compounds  they  have  found  between  the 
constituents  definite  relations,  and  have  often  regarded 
it  as  a  general  property  of  compounds  that  they  can  be 
found  only  in  these  unalterable  relations."  .  .  .  "This 
assumption  rests  merely  on  distinguishing  between  solution 
and  combination,  thereby  confusing  the  properties  which 
separate  the  two  with  affinity  which  effects  combination." 

As  is  well  known,  this  is  what  led  to  the  historical  dis- 
cussion between  Berthollet  and  Proust,  which  extended 
over  several  years.  Berthollet  held  the  view  that  the 
law  of  definite  proportions  is  not  true,  Proust  maintaining 
that  it  is.  Berthollet  was  wrong  and  Proust  was  right  — 
the  laws  of  definite  and  multiple  proportions  are  funda- 
mental laws  of  chemistry. 

Berthollet  summarizes  his  conception  of  solution  as  fol- 
lows: "The  solvent  represents  a  degree  of  chemical  activity 
which  differs  only  in  degree  from  that  which  produces 
the  most  stable  compound."  Solution  is  thus  a  purely 
chemical  action  between  solvent  and  dissolved  substance, 
resulting  in  the  formation  of  a  real  chemical  compound. 

It  is  interesting  to  note,  that  of  the  two  leading  French 
chemists  of  the  beginning  of  the  nineteenth  century,  the 
one,  Lavoisier,  held  a  purely  mechanical  conception  of 
solution;  the  other,  Berthollet,  a  purely  chemical. 

1  See  Walden:  Die  Losungstheorien  in  Hirer  geschichtlichen  Aufeinander- 
Jolge,  p.  53. 


EARLIER  VIEWS  OK  THE  NATURE  OF  SOLUTION        25 

After  it  was  shown  that  the  laws  of  definite  and  mul- 
tiple proportions  really  hold,  it  seemed  that  we  must 
make  some  distinction  between  chemical  combination, 
which  always  takes  place  hi  certain  definite  proportions, 
and  solution  which  occurs  hi  any  proportion,  at  least  up 
to  a  certain  limit.  This  was  done  by  Gay-Lussac.  He 
did  so  by  assuming  that  the  force  which  leads  to  chemical 
combination  is  more  powerful  than  that  which  produces 
solution.  Still  we  have  a  purely  chemical  theory  of 
solution. 

Thomson.  —  The  laws  of  definite  and  multiple  propor- 
tions had  been  discovered,  and  Dalton  had  proposed  his 
atomic  theory.  This  was  clearly  enunciated  by  Thomson 
hi  his  frequently  cited  "System  of  Chemistry."  Thomson 
also  regarded  solution  as  a  chemical  process,  but  dis- 
tinguished between  two  kinds  of  solutions:  Solids  com- 
bine with  some  of  the  solvent  and  remain  solid  —  form 
hydrates;  solids  dissolve  in  liquids  becoming  themselves 
Jiquid.  The  latter  systems  are  the  more  strictly  termed 
solutions. 

In  the  first  decades  of  the  nineteenth  century,  men  of 
science  took  up  the  quantitative  study  of  solutions.  They 
determined  what  substances  are  soluble  in  various  solvents, 
and  the  degree  of  their  solubility.  They  investigated  the 
physical  properties  of  solutions,  and  then  turned  to  the 
study  of  the  behavior  of  solutions  when  various  forms  of 
energy  were  passed  through  them.  They  measured  the 
conductivity  of  solutions  for  heat,  for  light,  and  for  elec- 
tricity. 

Grotthuss.  —  A  name  well  known  in  the  development 
of  the  theories  of  electrolysis  is  that  of  Grotthuss.  He 
regarded  chemical  phenomena  as  the  results  of  the  phenom- 
ena of  galvanic  electricity.1  "A  galvanism  manifests  itself 
in  a  liquid  due  to  the  heterogeneous  elementary  parts  con- 
tained in  it.  No  chemical  action  manifests  itself  when  the 
electrical  forces  are  in  equilibrium.  When  the  equilibrium 

1  See  Ostwald's  Klasvikern,  No.  152. 


26  THE  NATURE  OF  SOLUTION 

is  destroyed  chemical  action  takes  place."  Grotthuss 
recognized  that  elementary  substances  can  be  liquid,  and 
the  parts  of  these  were  supposed  not  to  be  homogeneous. 
Here  was  an  apparent  discrepancy.  Grotthuss  assumed 
that  since  elementary  substances  which  apparently  con- 
tain no  heterogeneous  parts  are  liquid,  the  elements  need 
not  be  regarded  as  made  up  of  simple  atoms,  but  the 
metals  can  be  regarded  as  containing  at  least  positively 
and  negatively  charged  parts,  which  are  chemically  com- 
bined with  one  another. 

The  view  suggested  by  Grotthuss,  in  1808,  is  very  in- 
teresting in  the  light  of  the  recent  work  of  Thomson.1 
He  has  electrolyzed  hydrogen  and  has  shown  that  it  con- 
tains positively  charged  hydrogen  atoms  and  negatively 
charged  hydrogen  atoms.  Since  the  hydrogen  molecules 
are  neutral,  they  are  probably  composed  each  of  a  posi- 
tive atom  and  a  negative  atom  which,  when  united,  would 
give  a  neutral  molecule.  Thus,  the  result  found  by  Thom- 
son was  predicted  nearly  a  century  before  it  was  experi- 
mentally established.  In  reference  to  solution  proper 
Grotthuss  expresses  himself  thus:  "The  solution  of  a 
salt  in  water  appears  to  be  only  an  interpolation  of  its 
own  elementary  parts  in  the  active  galvanic  molecular 
circle  of  water,  as  the  elements  of  the  salt  contribute  to 
the  galvanic  molecular  activity  of  the  water.  Salts 
which  cannot  do  this  we  call  insoluble." 

Grotthuss  also  distinguished  between  "solution"  and 
"dissolution,"  but  used  these  terms  in  a  somewhat 
different  sense  from  that  with  which  we  have  become 
familiar.  "Solution  and  dissolution  are,  according  to  my 
views,  different  in  this;  the  former  is  the  conversion  by  a 
liquid  of  a  solid  into  a  liquid,  where  the  newly  formed 
liquid  product  is  not  separable  by  galvanic  electricity 
into  the  original  liquid  and  solid,  (e.g.,  salt  dissolved  in 
water) ;  the  latter,  on  the  contrary,  is  the  conversion  by  a 
liquid  of  a  solid  into  a  liquid,  the  new  liquid  product 

1  Nature,  62,  451  (1895). 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        27 

being  decomposable  by  galvanic  electricity  into  its  ele- 
ments (e.g.,  acids  and  bases)." 

We  must  not  confuse  these  views  of  Grotthuss,  expressed 
in  1818,  with  his  ideas  announced  about  ten  years  earlier, 
which  are  the  basis  of  his  well-known  theory  of  electrolysis. 
He  earlier  maintained  that  hi  a  liquid  like  water  the 
hydrogen  atoms  and  oxygen  atoms  are  firmly  and  fixedly 
united.  A  hydrogen  atom  once  combined  with  a  given 
oxygen  atom  is  always  combined  with  that  same  oxygen 
atom.  It  becomes  separated  from  it  only  when  the  cur- 
rent is  passed  through  the  solution.  More  of  this  when 
the  theories  of  electrolysis  are  discussed.  His  later  views 
which  we  are  now  considering  are  almost  exactly  the 
opposite  of  this.  The  dissolved  particles  are  continually 
changing  partners.  In  a  salt  like  potassium  bromide,  the 
potassium  which  at  one  moment  is  combined  with  any 
given  bromine  atom,  the  next  moment  is  combined  with  a 
different  bromine  atom,  this  exchange  of  parts  going  on  con- 
tinually. The  relation  of  these  conceptions  to  the  theory 
of  electrolytic  dissociation  will  become  apparent  when  we 
come  to  the  consideration  of  that  theory. 

Berzelius*  Electrochemical  Theory.  —  This  brings  us  to 
the  electrochemical  theory  of  the  great  Swedish  chemist, 
Berzelius.  His  electrochemical  theory  must  be  briefly 
discussed  in  order  that  we  may  understand  his  views  in 
reference  to  the  nature  of  solution. 

That  there  is  an  intimate  relation  between  electrical 
action  and  chemical  action  was  rapidly  being  recognized 
by  men  of  science.  The  primary  cell  had  been  discovered 
by  the  Galvanis  and  by  Volta.  In  this  cell  there  was 
apparently  chemical  action,  and  electricity  resulted. 

Davy  had  constructed  his  enormous  voltaic  pile, 
consisting  of  some  four  hundred  couples,  and  had  plunged 
the  poles  of  this  pile  apparently  into  about  everything 
which  he  could  find,  to  see  what  would  happen.  He  had 
studied  the  electrolysis  of  aqueous  solutions  of  acids, 
bases,  and  salts.  He  had  decomposed  fused  sodium  and 


28  THE  NATURE  OF  SOLUTION 

potassium  hydroxides  by  means  of  the  current,  and  had  ob- 
tained metallic  sodium  and  metallic  potassium. 

Chemical  action  produced  electricity.  The  resulting 
electricity  decomposed  chemical  compounds.  This  would 
indicate  that  there  is  some  close  relation  between  the 
force  that  holds  together  the  constituents  of  chemical 
compounds,  and  electricity. 

Berzelius,  impressed  with  this  idea,  developed  his  elec- 
trochemical theory.  All  atoms  are  charged  both  posi- 
tively and  negatively,  but  there  is  always  a  preponderance 
of  one  or  the  other  kind  of  electricity  upon  every  atom, 
so  that  we  can  say  that  any  given  atom  is  charged  either 
positively  or  negatively.  A  positive  atom  and  a  negative 
atom  are  drawn  together  by  the  electrostatic  attraction 
of  their  opposite  charges,  and  they  combine  and  form  a 
chemical  compound.  Chemical  action  in  terms  of  this 
theory  is,  then,  nothing  but  the  electrostatic  attraction  of 
oppositely  charged  parts. 

The  consequences  of  this  electrical  theory  of  chemical 
action  are  very  interesting.  If  chemical  combination  is 
nothing  but  the  result  of  electrostatic  attraction,  the 
properties  of  the  resulting  compounds  must  be  a  function 
of  the  nature  of  the  electrical  charges  on  the  atoms  in  the 
compound.  This  is  a  necessary  consequence  of  the  electro- 
chemical theory  of  Berzelius. 

Facts  were  soon  discovered  which  seemed  to  militate 
against  this  theory.  Trichloracetic  acid  has  properties 
which  resemble  closely  those  of  acetic  acid.  Trichlor- 
acetic acid  contains  three  chlorine  atoms  instead  of  three 
hydrogen  atoms,  and  the  chlorine  atoms  were  supposed  to 
be  charged  negatively.  If  this  were  a  fact  it  would  be 
difficult  to  explain  in  terms  of  the  Berzelius  theory.  The 
electrical  condition  of  the  compound  is  changed  without 
changing  its  general  properties.  It  is  well  known  that  J.  J. 
Thomson1  has  shown  comparatively  recently  that  the 

1  Nature,  52,  452  (1895)  and  Author's  Elements  of  Physical  Chemistry, 
4th  edition,  p.  362.  (The  Macmillan  Co.) 


EARLIER  VIEWS  OF  THE   NATURE  OF  SOLUTION         29 

chlorine  atoms  in  trichloracetic  acid  are  charged  positively 
like  the  hydrogen  atoms  which  they  replace,  and  the 
above-mentioned  fact  is  now  in  perfect  accord  with  the 
Berzelius  theory. 

We  recognize  today  that  the  Berzelius  theory  contains 
a  large  element  of  truth. 

Berzelius'  Theory  of  Solution.  —  Having  hi  mind 
Berzelius'  theory  of  chemical  action,  we  can  now  under- 
stand his  views  as  to  the  nature  of  solution.  Solution 
differs  from  chemical  action  hi  that  it  takes  place  without 
any  electrical  act.  A  dissolved  substance  is  more  active 
chemically  than  the  same  substance  when  not  in  solution, 
and  it  must,  therefore,  preserve  the  electrical  difference  of 
its  parts. 

Berzelius  regarded  solution  as  due  to  some  unknown 
modification  of  affinity,  which  gave  a  very  different  result 
from  the  action  of  affinity  itself.  Solution  he  regarded  as 
essentially  a  mechanical  process. 

"The  kind  of  force  which  effects  the  solution  of  a  solid 
in  a  liquid  is  not  identical  with  the  force  which  brings 
about  chemical  union,  and  is  not  to  be  confused  with  it. 
-When  the  latter  acts  heat  is  formed,  while  the  force 
which  produces  solution  absorbs  heat  and  lowers  the  tem- 
perature. "* 

This  is  a  rather  sharp  distinction,  and,  considering 
the  tune  when  it  was  suggested,  has  a  peculiar  interest  of 
its  own. 

Berzelius,  of  course,  knew  that  heat  is  often  evolved  in 
processes  of  solution.  He  accounted  for  this  as  due  to  the 
dissolved  substance;  and  this  evolution  of  heat  could 
readily  be  larger  than  the  heat  absorption  due  to  the 
process  of  solution.  Dissolved  substances  do  not  alter  their 
chemical  properties,  they  simply  alter  their  state  of  aggre- 
gation, and  become  chemically  more  active  due  to  the 
fact  that  they  are  dissolved.  One  fundamental  difference 
between  solutions  and  chemical  compounds  had  been  noted 

1  Lehrbuch  der  Chemie  I  (1842). 


30  THE  NATURE  OF  SOLUTION 

by  several,  including  Berzelius.  Chemical  compounds 
are  formed  in  terms  of  the  law  of  definite  proportions. 
Solution  takes  place  in  any  quantity,  up  to  a  certain 
maximum  limit.  This,  as  we  now  know,  is  a  funda- 
mental difference.  This  brings  us  to  the  views  of  Gay- 
Lussac. 

Gay-Lussac.  —  The  role  played  by  Gay-Lussac  in  the 
development  of  our  knowledge  of  solution  is  historically 
not  only  of  interest,  but  very  important.  He  seems  to 
have  been  the  first  to  recognize  clearly  a  relation  between 
solutions  and  gases.  In  his  own  words,1  "Like  the  elas- 
ticity of  vapors,  the  solution  of  a  substance  varies  with 
the  temperature.  It  is  without  doubt  also  dependent  on 
the  reciprocal  affinity  of  solute  and  solvent;  but  since  the 
effects  of  affinity  are  not  variable  with  temperature,  while 
those  of  solution  depend  fundamentally  upon  it,  it  would 
be  difficult  not  to  admit  that  in  solution,  as  in  vaporiza- 
tion, the  product  is  essentially  limited  for  each  degree  of 
temperature  by  the  number  of  molecules  which  can  exist 
in  a  given  portion  of  the  solvent.  For  the  same  reason 
that  the  elastic  molecules  are  precipitated  by  lowering  the 
temperature,  they  are  separated,  and  probably  also,  like 
the  latter,  by  the  compression  and  reduction  of  the  volume 
of  the  solvent." 

We  now  come  to  the  important  point  hi  the  paper  by 
Gay-Lussac.2 

"  Solution  is,  then,  fundamentally  related  to  vaporisa- 
tion in  the  sense  that  both  are  dependent  on  temperature 
and  obey  its  variations.  Therefore,  there  ought  to  be 
between  the  two,  if  not  a  complete  identity  of  result,  at 
least  a  close  analogy;  their  essential  difference  consisting 
hi  this,  that  the  gaseous  molecules  do  not  need  a  solvent 
to  maintain  them  in  a  given  space.  Their  repulsive 
force  is  sufficient  for  this  purpose.  On  the  other  hand,  in 
a  solution  of  a  solid  or  a  liquid,  the  molecules  are  not  able 

1  Ann.  Chim.  Phys.  [2],  70,  424  (1839). 

2  Ibid.,  425. 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        31 

to  maintain  themselves  in  the  space,  if  they  are  not  united 
by  affinity  to  the  molecules  of  the  solvent. 

"The  analogies  between  solution  and  vaporization 
extend  also  to  the  effects  on  them  of  variation  in  tempera- 
ture; and  as  it  appears  to  me  unquestionable  that  the 
elastic  force  of  the  vapor  of  a  substance  is  independent 
of  the  state  of  this  substance  and  of  the  cohesion  of  its 
molecules,  since  the  one  remains  constant  while  the  other 
varies,  from  analogy  I  am  inclined  to  admit  that  solu- 
tion is  equally  independent  of  cohesion. 

"  However,  admitting  the  analogies  between  vaporiza- 
tion and  solution,  we  must  ask  why  it  is,  that,  while  the 
elastic  force  of  vapors  obeys  a  regular  law,  increasing 
regularly  with  the  temperature,  the  solubility  of  certain 
salts  shows  a  maximum  and  then  decreases." 

This,  Gay-Lussac  explains  as  due  to  a  change  in  the 
composition  of  the  compound  at  the  temperature  in 
question. 

These  views  of  Gay-Lussac  are  very  important,  in  that 
they  are  the  forerunner  of  quantitative  relations  between 
solutions  and  gases,  pointed  out  nearly  a  half  century 
Jater  by  Van't  Hoff,  which,  as  we  shall  see,  are  by  far  the 
most  important  contributions  to  our  knowledge  of  solu- 
tions that  have  thus  far  been  made. 

Gay-Lussac  first  called  attention  to  the  analogy  between 
matter  in  the  gaseous  and  in  the  dissolved  states. 

There  were  quite  a  number  of  prominent  men  of  science 
of  this  period  who  advocated  a  chemical  theory  of  solu- 
tion, i.e.,  that  solution  is  due  to  a  force  acting  between 
the  solvent  and  the  dissolved  substance  which  caused  the 
two  to  combine. 

Williamson.  —  The  English  chemist,  Williamson,  pub- 
lished a  paper  in  1851,  "On  the  Theory  of  Ether  Forma- 
tion,"1 which  has  proved  to  have  great  value  in  more 
directions  than  one.  Its  greatest  service  has  probably 
been  in  pointing  out  the  real  nature  of  chemical  equi- 

1  Lieb.  Ann.,  77,  37  (1851). 


32  THE  NATURE  OF  SOLUTION 

librium  —  that  it  is  not  a  statical  condition  as  had  been 
hitherto  supposed,  but  a  dynamical  one,  representing  that 
condition  where  the  two  opposite  reactions  have  equal 
velocities. 

This  paper  is  also  of  interest  and  importance  in  its 
bearing  on  the  development  of  the  theory  of  solution,  and 
in  this  connection  must  now  be  considered.  Williamson 
pictured  the  action  of  sulphuric  acid  on  ethyl  alcohol  as 
follows: 

H  H 

S04  S04 


This  means  that  when  sulphuric  acid  reacts  with  ethyl 
alcohol,  there  are  formed  ethyl  sulphuric  acid  and  water. 

When,  however,  more  ethyl  alcohol  is  allowed  to  re- 
act with  ethyl  sulphuric  acid,  the  reaction,  according  to 
Williamson,  takes  place  in  the  sense  of  the  following 

equation: 

H  H 

SO4        S04 
C2H8  H 

H        "  C2H6 
O  O 

C2H6 


In  this  part  of  the  reaction  sulphuric  acid  and  ether  are 
formed. 

In  the  first  part  of  the  reaction,  as  will  be  seen,  the 
hydrogen  of  the  sulphuric  acid  is  replaced  by  the  ethyl 
group.  In  the  second  stage  the  ethyl  group  of  the 
ethyl  sulphuric  acid  is  replaced  by  hydrogen  —  the  one 
reaction  being,  in  a  sense,  just  the  reverse  of  the  other. 

Williamson  would  explain  these  and  similar  facts  in 
the  following  manner,  to  quote  his  own  words:1  "We  are 
thus  led  to  the  assumption  that  in  a  number  of  molecules 
of  any  compound,  a  continual  exchange  is  going  on 
between  the  elements  in  the  compound.  Let  us  take,  for 

1  Lieb.  Ann.,  77,  46  (1851). 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        33 

example,  a  vessel  containing  hydrochloric  acid;  it  is  filled 
with  a  large  number  of  molecules  of  HC1.  The  above 
considerations  would  lead  us  to  the  assumption  that  an 
atom  of  hydrogen  does  not  remain  quietly  attached  to 
the  atom  of  chlorine  with  which  it  was  first  combined, 
but  there  is  going  on  a  continued  exchange  of  place  with 
other  hydrogen  atoms.  This  change  is  for  us,  of  course, 
not  directly  detectable,  because  one  atom  of  hydrochloric 
acid  is  just  like  another." 

This  view,  that  hi  solution  there  is  a  constant  niter- 
change  of  parts,  implies  that  in  solution  things  are  more 
or  less  broken  down  into  their  parts.  Indeed,  Williamson 
went  farther  than  was  necessitated  by  the  facts,  and 
assumed  more  decomposition  hi  solution  than  was  neces- 
sary, as  was  shown  by  subsequent  work. 

Clausius.  —  The  physicist  Clausius,  in  1856,  published 
a  paper,1  "On  the  Conduction  of  Electricity  hi  Electro- 
lytes," which  proved  to  be  of  very  great  importance.  In 
this  paper  Clausius  criticizes  Williamson  for  having  gone 
too  far,  indeed,  much  farther  than  the  facts  required  or 
justified.2  "  Williamson  appears  to  have  assumed  a 
,  greater  interchangeability  in  the  grouping  of  the  part 
molecules  than  is  necessary  to  explain  the  conduction  of 
electricity.  He  speaks  of  one  hydrogen  atom  continually 
changing  places  with  another,  while  to  account  for  the 
conduction  of  electricity,  it  is  sufficient  to  assume  that 
as  the  whole  molecules  strike  against  one  another,  there 
occurs,  perhaps  relatively  seldom,  an  exchange  of  the  part 
molecules." 

Clausius  then  proceeded  to  develop  his  own  theory.  If 
we  assume  that  an  electric  current  is  necessary  to  decom- 
pose the  whole  molecules  into  part  molecules,  then  the 
force  necessary  to  break  down  these  whole  molecules 
must  have  a  certain  value.  If  this  force  were  not  applied, 
i.e.,  if  the  force  were  smaller  than  that  necessary  to  decom- 
pose the  molecules,  no  current  would  pass  and  no  elec- 

1  Pogg.  Ann.,  101,  338  (1856).  2  Ibid.,  353. 


34  THE  NATURE  OF  SOLUTION 

trolysis  would  take  place.  As  the  force  is  increased  until 
it  reaches  the  decomposition  value,  there  would  be  no 
decomposition  until  this  value  is  reached;  then,  sud- 
denly a  large  number  of  molecules  would  be  decomposed 
by  the  current. 

The  facts  are  directly  opposed  to  this  conclusion.  It 
had  been  shown1  that  an  infinitesimal  force  is  capable 
of  effecting  a  corresponding  electrolysis  of  water;  and 
the  amount  of  the  electrolysis  is  proportional  to  the  force 
acting  upon  it.  A  force  altogether  too  small  to  decom- 
pose even  one  molecule  of  water  can  electrolyze  water. 
The  conclusion  from  this  fact  is  obvious.  If  a  force  too 
weak  to  decompose  water  can  electrolyze  it,  the  current 
must  find  the  water  already  decomposed,  at  least  to  some 
extent. 

Clausius'  theory  of  the  nature  of  solutions  of  electro- 
lytes is,  then,  as  follows:  Such  a  solution  consists  chiefly 
of  "  whole  molecules,"  as  they  were  termed,  but  there 
are  also  present  a  few  "part  molecules."  The  whole 
molecules  are  breaking  down  and  the  part  molecules 
recombining  all  the  tune. 

When  an  electromotive  force  is  impressed  upon  such  a 
solution,  the  current  does  not  have  to  break  down  the 
molecules  of  the  electrolyte.  It  finds  some  of  them  (part 
molecules)  already  in  a  state  of  decomposition.  The 
current  exerts  only  a  directing  action,  driving  the  posi- 
tively charged  parts  in  the  direction  of  the  negative  pole, 
and  the  negatively  charged  parts  in  the  direction  of  the 
positive  pole. 

We  can  now  understand  in  terms  of  the  theory  of 
Clausius,  how  a  current  too  weak  to  decompose  even  one 
molecule  of  water  can  effect  the  electrolysis  of  water.  It 
is  not  necessary  to  decompose  the  water.  The  current 
finds  the  water  already  slightly  decomposed  into  "part 
molecules."  It  is  only  necessary  for  the  current  to  exert 

1  See  Author's  Elements  of  Physical  Chemistry.     4th  edition,  p.  368  (Mac- 
millan  Co.)- 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        35 

a  directing  action  on  these  parts,  and  electrolysis  is  the 
result.  The  historical  importance  of  the  theory  of  Clausius 
will  be  seen  when  we  come  to  consider  the  theory  of 
electrolytic  dissociation. 

Kopp.  —  Kopp  distinguished  between  compounds  which 
obey  the  law  of  definite  proportions,  and  those  hi  which 
the  constituents  are  present  hi  varying  proportion.  The 
latter  are,  in  general,  solutions.  Kopp  defines  a  solution 
of  a  salt  in  water  as,  "a  compound,  the  constituents  of 
which  are  present  in  varying  quantities."  A  concentrated 
solution  is  an  unstable  compound  of  the  dissolved  substance 
with  water.  The  composition  of  this  compound  can  be 
easily  varied  by  change  in  temperature,  and  by  many  other 
means. 

Kopp  points  out  one  marked  difference  between  the 
properties  of  stable  or  well-defined  chemical  compounds 
and  solutions.  When  substances  unite  and  form  a  defi- 
nite, stable  compound,  most  of  the  properties  of  the  com- 
pound are  very  different  from  those  of  its  constituents. 
Indeed,  we  know  today  that  about  the  only  property  of 
the  constituents  which  persists  hi  chemical  action  is  mass. 
-To  within  the  limit  of  accuracy  of  our  most  refined  chemi- 
cal balance  there  is  no  change  in  mass  in  chemical  reac- 
tion.1 

The  properties  of  solutions,  on  the  other  hand,  are  often 
additive,  if  they  are  the  sum  of  the  properties  of  the  con- 
stituents. The  properties  of  other  solutions  are  the  mean 
of  the  properties  of  the  several  constituents.  In  still 
other  solutions  the  properties  of  one  constituent  affect  those 
of  the  other  constituents  to  a  greater  or  less  extent.  In 
these  cases  the  properties  of  the  solution  lie  somewhere 
between  those  of  the  several  constituents.  This  is  an 
important  distinction  between  true,  stable,  chemical  com- 
pounds, and  the  more  unstable  hydrates  which,  as  we  shall 
see,  exist  so  often  in  aqueous  solution. 

Guldberg  and  Waage.  —  The  landmarks  in  the  history 
1  Landolt:  Zeit.  phys.  Chem.,  12,  1  (1893);  65,  589  (1906). 


36  THE  NATURE  OF  SOLUTION 

of  any  branch  of  science  are  the  generalizations  or  laws 
that  have  been  reached  correlating  the  phenomena  in 
question.  The  highest  aim  of  the  man  of  science  is  to 
discover  such  a  law. 

Two  names  which  will  always  live  in  the  history  of 
chemistry,  because  the  men  discovered  not  only  a  law  of 
chemical  action,  but  one  of  the  most  wide  reaching  and 
important  laws  of  all  chemistry,  viz.,  the  law  of  mass  ac- 
tion, are  those  of  the  Norwegian  physicist,  Guldberg,  and 
the  Norwegian  chemist,  Waage. 

These  men  have  a  place  also  hi  the  history  of  the 
development  of  our  present  conception  of  the  nature  of 
solution.  They  recognized  in  general  the  two  classes  of 
chemical  compounds,  those  in  which  the  constituents  are 
present  in  definite  relations,  and  those  in  which  they  are 
not.  They  also  recognized  the  existence  of  compounds 
intermediate  between  the  above  two  classes.  These  are 
present  in  solution.  Indeed,  there  are  all  gradations 
from  compounds  obeying  rigidly  the  law  of  definite  propor- 
tions to  those  which  do  not  conform  at  all  to  this  law. 
All  of  these  substances  are  produced  by  the  action  of  the 
same  force. 

The  distinctively  new  feature  introduced  by  Guldberg 
and  Waage,  in  connection  with  solution,  is  that  hi  solution 
the  molecules  of  the  dissolved  substance  are  not  free  from 
one  another,  but  they  are  united  hi  larger  or  smaller 
groups,  and  the  size  of  these  groups  varies  with  the  temper- 
ature of  the  solution,  and  probably  also  with  the  dilution. 

The  "groups"  of  Guldberg  and  Waage  are,  of  course, 
groups  of  the  same  kind  of  molecules.  They  are  what  we 
today  call  polymerized  molecules. 

Valson.  —  In  1870  a  very  important  contribution  to 
our  knowledge  of  solutions  was  made  by  Valson.1  From 
the  study  of  the  capillary  action  of  liquids,  Valson  dis- 
covered what  he  termed  the  "law  of  capillary  moduli. " 
This  he  describes  in  the  following  words.2  "Let  us  call  a 

1  Ann.  Chim.  Phys.  [4],  20,  361  (1870).  2  Ibid.,  364. 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        37 

salt  Mm:  M  being  the  metallic  radical,  and  m,  the  non- 
metallic.  If  now  we  pass  to  a  second  salt  Mm',  con- 
taining the  same  metal  united  with  a  different  non-metal, 
the  capillary  effect  due  to  the  radical  M  remains  constant 
whatever  the  nature  of  m.  Similarly,  if  we  pass  from  the 
salt  Mm  to  a  salt  M  'm,  the  capillary  effect  due  to  the  radi- 
cal m  will  be  the  same  whatever  the  nature  of  the  metal. 
Finally,  if  we  pass  from  the  salt  Mm  to  a  salt  M'm', 
in  which  both  radicals  have  at  the  same  tune  been 
changed,  the  total  effect  would  be  equal  to  the  sum  of  the 
effects  produced  by  the  two  radicals  taken  separately; 
provided  the  solutions  are  sufficiently  dilute  and  are 
moreover  hi  a  state  of  normal  solution,  i.e.,  that  they 
contain  one  equivalent  of  the  salt  dissolved  hi  one  liter 
of  water. 

"  I  give  to  this  effect  the  name  capillary  modulus,  and 
I  am  able  to  enunciate  the  following  law,  which  holds  for 
saline  solutions. 

1st.  "The  modulus  of  a  metallic  radical  is  constant 
and  independent  of  the  non-metallic  radical  with  which  it 
is  associated. 

2nd.  "The  modulus  of  a  non-metallic  radical  is  constant, 
and  independent  of  the  metallic  radical  with  which  it  is 
associated. 

"If  the  two  radicals  change  at  the  same  time,  the 
modulus  is  equal  to  the  sum  of  the  two  partial  moduli. 

"The  modulus  of  a  salt  is  equal  to  the  sum  of  the 
moduli  of  its  two  radicals. " 1 

Valson  found  that  these  relations  held  only  for  fairly 
dilute  solutions. 

"  The  determinations  of  the  moduli  were  made  in 
solutions  containing  one  equivalent  of  salt  dissolved  hi 
one  liter  of  water.  It  might  at  first  seem  better  to  use 
more  concentrated  solutions  containing,  for  example,  two 
equivalents  of  salt  "  .  .  .  "  On  comparing  dilute  solutions, 
experiment  shows  that  the  capillary  effects  are  approxi- 

1  Ann.  Chim.  Phys.  [4],  20,  386  (1870). 


38  THE  NATURE  OF  SOLUTION 

mately  proportional  to  the  quantities  of  the  substances 
used;  but  this  is  not  the  case  for  concentrated  solutions, 
where  the  actions  of  the  molecules  are  not  independent  of 
one  another.  It  thus  seems  that  in  order  that  the  saline 
molecules  may  be  regarded  as  having  acquired  a  state  of 
freedom,  it  is  necessary  that  they  be  in  a  sufficiently  di- 
lute solution." 1 

Thus  was  pointed  out  clearly  the  difference  between 
dilute  and  concentrated  solutions.  The  additive  proper- 
ties which  were  found  in  dilute  solutions  did  not  hold  in 
concentrated.  We  shall  see  that  this  distinction  between 
dilute  and  concentrated  solutions  is  one  of  fundamental 
importance,  the  meaning  of  which  has  been  only  com- 
paratively recently  pointed  out. 

Valson  showed  later  that  the  specific  gravities  of  solu- 
tions of  salts  conform  to  the  same  general  relations  as  their 
refractivities  —  for  dilute  solutions  they  are  additive. 

Favre  and  Valson.  —  Favre  and  Valson,2  in  their  paper 
on  the  relation  between  thermal  effects  and  contraction 
in  solution,  showed  that  what  they  termed  "moduli  of 
coersion"  were  analogous  to  moduli  of  thermal  effects, 
density  and  capillarity. 

"Do3  not  the  results  in  the  last  table  lead  us  to  ask 
whether  the  solvent  action  of  water  on  salts  is  not  to  dis- 
sociate them  into  their  elements,  and  to  place  them,  if 
not  hi  a  state  of  complete  freedom,  at  least  hi  a  state  of 
reciprocal  independence,  which  it  would  be  difficult  at 
present  to  define,  but  which  is  very  different  from  their 
original  state. 

"This  independence  of  the  saline  elements  on  standing 
is  not  simply  an  hypothesis.  It  is  a  conclusion  to  which 
we  have  repeatedly  been  led  in  the  course  of  our  researches. 
Without  going  into  details  we  shall  content  ourselves  with 
recalling  what  we  have  said  in  general  on  thermoneu- 

i  Ann.  Chim.  Phys.  [4],  20,  386,  (1870). 

*  Compt.  Rend.,  76,  330,  798,  925,  1000  (1872). 

3  Ibid.,  pp.  1004  and  1005. 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION         39 

trality,  on  moduli  of  density  and  capillary  action,  and 
finally,  on  the  moduli  of  coersion.  ...  In  sufficiently 
dilute  solutions  of  salts,  each  of  the  metallic  or  non- 
metallic  elements  of  the  salts  always  produces  the  same 
effects,  and  these  are  independent  of  the  other  elements 
which  are  present." 

Favre  and  Valson  then  draw  the  following  conclusion: 

"Solution  gives  to  the  elements  of  the  dissolved 
substance  a  reciprocal  independence,  and  the  inner  mechan- 
ical work  necessary  to  produce  this  effect  is  measured  by 
the  changes  in  volume  which  accompany  solution  and, 
consequently,  by  the  quantity  of  heat  set  free  when  the 
same  effects  of  coersion  are  produced  directly  on  the 
solvent  liquid." 

When  we  come  to  consider  the  theory  of  electrolytic 
dissociation,  we  shall  see  that  the  work  and  conclusions 
of  Favre  and  Valson  bear  directly  upon  it.  Had  they 
regarded  the  salt  parts  as  charged  when  dissociated  from 
one  another,  and  had  they  pointed  out  a  method  for 
measuring  the  magnitude  of  the  dissociation,  they  would 
have  approached  very  closely  to  the  theory  of  electrolytic 
dissociation. 

As  we  shall  see,  however,  the  discovery  of  this  epoch- 
making  generalization  was  left  for  another. 

Landolt.  —  A  large  amount  of  work  was  done  at  this 
time  which  brought  out  the  additive  properties  of  aqueous 
solutions  of  salts.  Landolt l  studied  the  power  of  substances 
to  rotate  the  beam  of  polarized  light,  and  found  for  such 
substances  as  tartaric  acid  and  its  salts  that  the  rotation 
is  independent  of  the  optically  inactive  constituent  pres- 
ent. This  would  indicate  that  hi  solution  the  molecule 
is  broken  down  into  its  constituents. 

Gladstone.  —  From  his  work  on  the  refractivity  of 
solutions,  Gladstone2  had  earlier  shown  that  the  refractive 
power  of  a  salt  is  the  sum  of  two  constants,  the  one 

*  Eer.  d.  chem.  GeselL,  6,  1073  (1873). 
»  Phil  Mag.,  36,  313  (1868). 


40  THE  NATURE  OF  SOLUTION 

depending  for  its  value  upon  the  metal,  and  the  other 
upon  the  non-metallic  constituent  present. 

Kohlrausch.  —  Friedrich  Kohlrausch1  had  worked  out 
his  well-known  method  for  measuring  the  conductivity 
of  solutions  of  acids,  bases,  and  salts,  and  had  studied 
the  power  of  these  solutions  to  conduct  the  electric  cur- 
rent. He  had  discovered  the  law,  which  bears  his  name, 
of  the  independent  migration  velocity  of  the  ions.  This 
says  that  the  conductivity  of  any  solution,  referred  to 
molecular  quantities  of  the  dissolved  substance  —  the 
molecular  conductivity  —  is  the  sum  of  two  constants, 
the  one  depending  for  its  value  on  the  cation,  the  other 
on  the  anion. 

These  and  many  other  results  which  were  obtained 
about  this  time  from  the  investigation  of  the  physical 
properties  of  aqueous  solutions,  all  pointed  in  one  direc- 
tion, viz.,  that  the  properties  of  such  solutions  are  the 
sums  of  two  values,  one  depending  on  the  metallic  con- 
stituent of  the  salt,  the  other  on  the  non-metallic  con- 
stituent. The  simplest  explanation  of  this  fact  would  be 
that  hi  such  solutions  a  salt  is  broken  down  into  two 
constituents,  the  one  metallic,  the  other  non-metallic,  and 
that  each  of  these  constituents  has  its  own  distinctive 
and  characteristic  properties  which  are  not  affected  by 
the  presence  of  the  other.  This  is  the  simplest  and  most 
obvious  interpretation  of  the  facts.  We  shall  see  later 
whether  it  is  true. 

Berthelot.  —  It  was  about  1860  that  two  men,  inde- 
pendently, took  up  the  study  of  the  amounts  of  heat  that 
are  liberated  in  chemical  reactions.  The  object  of  this 
work  was  not  simply  to  make  thermochemical  measure- 
ments. Far  from  it.  It  was  recognized  that  the  cause 
of  all  chemical  reaction  is  to  be  found  in  the  thermal  change 
that  always  takes  place  whenever  substances  react  chemi- 
cally, and  which  was  referred  to  up  to  that  time  as  the 
thermal  change  that  "accompanies  chemical  reactions." 

1  Wied.  Ann.,  6,  168  (1879). 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION         41 

A  moment's  thought  will  show  that  this  was  confusing 
cause  and  effect.  It  was  the  energy  change  that  brought 
about  the  material  change.  The  study  of  the  former 
was,  then,  far  more  fundamental  than  the  study  of  the 
latter,  and  absolutely  essential,  if  we  would  ever  trans- 
form chemistry  from  pure  empiricism  into  science.1 
This  was  recognized  by  Thomsen  in  Copenhagen,  and  es- 
pecially by  Berthelot  in  Paris. 

The  thermochemical  investigations  of  Berthelot  were 
published  in  two  large  volumes2  in  1879.  This  work 
can  be  referred  to  here  only  in  so  far  as  it  deals  with 
the  nature  of  solution.  "These3  inequalities  between  the 
specific  heats  of  solutions  of  salts,  and  those  of  their  com- 
ponents, water  and  anhydrous  salt,  appear  to  be  due," 
according  to  Berthelot,  "to  the  formation  right  in  the 
solution  of  certain  definite  hydrates,  comparable  with  the 
saline,  crystallized  hydrates.  There  is,  however,  this  dif- 
ference, that  the  dissolved  hydrates  exist  most  frequently 
in  the  liquid  hi  a  state  of  partial  dissociation,  the  amount 
of  the  dissociation  varying  with  the  quantity  of  the  water, 
and  with  the  temperature. 

"The4  phenomena  presented  by  a  normal  solution  are, 
in  a  certain  sense,  intermediate  between  those  of  the  simple 
mixture  and  the  true  compound.  Indeed,  the  tendency  to 
unite  and  form  a  homogeneous  system  indicates  a  real 
affinity  between  the  solid  and  the  solvent.  On  the  other 
hand,  this  union  ceases  under  the  influence  of  simple 
evaporation,  and  it  takes  place  apparently  in  proportions 
which  vary  continuously  with  the  temperature."  .  .  .  "It5 
appears  to  me  probable,  however,  that  the  distinctive  fea- 
ture of  solution,  in  the  true  sense  of  the  term,  lies  in  the 
formation  of  certain  definite  compounds  between  the  sol- 
vent and  the  dissolved  substance.  These  are  the  definite 

1  See:  A  New  Era  in  Chemistry,  by  the  Author  (D.  Van  Nostrand  Co,  N.  Y.)- 

2  Essai  de  Mecanique  Chimique,  fondee  sur  la  Thermochemie  (1879). 
8  Essai  de  Mecanique  Chimique,  vol.  I,  p.  507. 

«  Ibid.,  vol.  n,  p.  160. 
•  Ibid.,  p.  161. 


42  THE  NATURE  OF  SOLUTION 

hydrates  that  are  formed  right  in  the  solution  itself, 
between  the  salt  and  the  water  that  is  present  in  the 
solution,  hydrates  analogous  to,  or  identical  with,  those 
definite  hydrates  of  the  same  compounds  known  in  the 
crystalline  condition. 

"Let  us  insist  on  this  point  which  is  very  important  in 
chemical  statics. 

"The1  chemical  statics  of  solutions  depends  on  the 
actual  state  of  these  different  compounds,  hydrates  or 
anhydrous  substances,  actually  existing  hi  the  solutions. 
It  is  therefore  extremely  important  to  be  able  to  define  this 
state." 

We  find  also  on  page  163  the  following:  "What- 
ever be  the  hypothesis  relative  to  the  state  of  the  dis- 
solved substances,  the  actual  existence  in  solution  of 
certain  definite  hydrates  formed  by  acids,  alkalies,  and 
salts  is  independent  of  it.  It  can  be  established  by  many 
demonstrations,  on  the  one  hand  from  the  physical  proper- 
ties, on  the  other  from  the  chemical  properties  of  solu- 
tions." 

Berthelot  then  gives  a  number  of  these  demonstrations, 
but  it  would  lead  us  too  far  to  take  them  up  here.  What 
has  already  been  stated  will  suffice  to  make  clear  the 
views  of  Berthelot  on  the  nature  of  solution. 

Thomsen.  —  The  other  prominent  thermochemist  was 
the  Dane,  Julius  Thomsen.  He  published  the  results  of 
his  thermochemical  investigations  in  his  four-volume 
work,  "  Thermochemische  Untersuchungen."  His  ideas 
on  solutions  are  interesting  mainly  in  connection  with  the 
relations  between  solutions  and  gases.  The  closing  para- 
graph2 of  the  first  volume  of  his  great  work  runs  thus: 
"  These  investigations  lead  to  the  conclusion,  that  many 
apparent  irregularities  in  the  heat  tone,  which  have  been 
observed  when  the  substances  resulting  from  the  reaction 
of  acids  and  bases  are  not  all  soluble,  would  disappear  if 

i  Essai  de  Mecanique  Chimique,  Vol.  I,  p.  163. 

*  Tbermochemische  Untersuchungen,  vol.  I,  p.  449  (1882). 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        43 

all  the  substances  both  before  and  after  the  reaction  were 
present  as  aqueous  solutions;  and  that  aqueous  solutions 
of  substances  contain  them  in  a  condition,  hi  which,  as  hi 
the  gaseous  state,  the  physical  properties  are  the  simplest 
possible,  and  can  be  compared  directly  with  one  another. " 

Thomsen  was  the  second  investigator  to  recognize  the 
relation  between  solutions  and  g'ases.  Gay-Lussac,  it  will  be 
recalled,  was  the  first.  The  importance  and  significance 
of  the  relation  will  be  seen  later. 

Raoult.  —  An  interesting  figure  in  chemistry  at  the 
tune  that  we  are  now  considering  was  the  Frenchman, 
F.  M.  Raoult.  He  carried  out  his  investigations  at  the 
University  of  Grenoble.  His  work  on  the  lowering  of 
the  freezing-point  and  the  lowering  of  the  vapor-tension  of 
solvents  by  dissolved  substances  has  become  a  classic. 

From  his  paper1  published  hi  1885  the  following  para- 
graphs are  translated.  "The2  preceding  relations  show 
that  in  all  the  salts  which  it  enters,  every  radical  plays 
nearly  the  same  r61e  and  has  the  same  constitution;  and 
that  it  always  produces  the  same  lowering  of  the  freezing- 
point,  which  is  nearly  independent  of  the  number  and  the 
nature  of  the  other  radicals  with  which  it  can  combine." 
Again,  "The3  molecular  lowering  of  the  freezing-point, 
produced  by  salts  formed  by  strong  monobasic  and  dibasic 
acids,  is  approximately  the  sum  of  the  partial  molecular 
lowerings  of  their  electropositive  and  electronegative  rad- 
icals." 

And  near  the  end  of  his  paper,  "The4  diminution  in 
capillary  height,  the  increase  in  density,  the  contraction  of 
protoplasm  (de  Vries'  work  on  osmotic  pressure),  the 
lowering  of  the  freezing-point,  in  brief,  most  of  the  physical 
effects  produced  by  salts  on  solvent  water  are  the  sum  of 
the  effects  produced  separately  by  the  electropositive  and 
electronegative  radicals  which  are  contained  in  them,  and 
which  act  as  if  they  were  simply  mixtures  hi  the  liquid." 

1  Ann.  Chim.  Phys.  [6],  4,  401  (1885).  2  Ibid.,  414. 

»  Ibid.,  417.  *  Ibid.,  427. 


44  THE  NATURE  OF  SOLUTION 

"This1  shows  that  salts  in  solution  in  water  ought  to  be 
regarded  as  systems  of  particles,  every  one  of  which  is 
formed  of  solitary  atoms,  and,  notwithstanding  its  condition 
of  combination  with  the  others,  largely  preserves  its  own 
individuality,  its  action,  and  its  characteristic  properties. 

"The2  special  constitution  which  we  are  compelled  to 
recognize  for  true  salts  dissolved  in  water,  does  not  hold  for 
other  substances,  nor  to  any  appreciable  extent  in  solvents 
other  than  water.".  .  .  "The3  metallic  salts  themselves,  in 
solvents  other  than  water,  show  nothing  peculiar,  and 
produce  the  same  molecular  lowering  as  all  other  sub- 
stances." 

Raoult  concludes  his  paper4  thus:  "From  this  we 
must  conclude  that  the  subdivision  of  the  saline  mole- 
cules into  electropositive  and  electronegative  radicals 
takes  place  in  fact  only  in  aqueous  solutions;  and  con- 
sequently, this  is  the  result  of  a  peculiar  chemical  act 
exerted  by  water  on  salts  dissolved  in  it." 

We  are  thus  coming  nearer  and  nearer  to  the  theory  of 
electrolytic  dissociation.  Valson  concluded  that  the  par- 
ticles in  salts  must  be  regarded  as  separated  from  one 
another.  Raoult  goes  one  step  farther,  and  concludes 
that  these  particles  must  be  regarded  the  one  as  electro- 
positive and  the  other  as  electronegative. 

Mendeleeff.  —  We  have  seen  that  a  large  number 
had  held  the  view  that  in  solution  there  is  some  kind  of 
combination  between  the  solvent  and  the  dissolved  sub- 
stance. Most  of  the  ideas  that  had  been  expressed  were 
vague  and  indefinite;  indeed  they  were  so  general  that 
it  was  difficult  to  test  them.  The  views  of  the  great 
Russian  chemist,  Mendeleeff,  while  somewhat  indefinite, 
are  far  more  concrete  than  those  of  his  predecessors;  and 
the  earlier  so-called  hydrate  theory  of  solution  bears 
his  name.  His  views  are  expressed  in  various  places  in 
his  "Principles  of  Chemistry." 

1  Ann.  Chim.  Phys.  [6],  4,  428  (1885). 

2  Ibid.,  428.          *  Ibid.,  429.  «"  Ibid.,  430. 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        45 

Reactions  in  aqueous  solution  depend  both  qualita- 
tively and  quantitatively  on  the  mass  of  the  water  present, 
in  a  word,  on  the  dilution  of  the  solution.  Many  sub- 
stances form  with  water  a  number  of  compounds,  e.g., 
many  substances  crystallize  with  different  amounts  of 
water,  and  each  one  of  these  substances  has  its  own  definite 
and  characteristic  properties. 

He  regarded  solutions  as  unstable,  but  definite,  chemical 
compounds  which  were  undergoing  dissociation.  The  num- 
ber of  hydrates  formed  by  any  one  substance  with  water 
was  small.  In  a  great  majority  of  cases  there  is  only 
one  hydrate  of  the  substance  hi  question. 

Mendeleeff  regarded  compounds  of  definite  composition 
as  only  a  special  case  of  compounds  of  indefinite  composi- 
tion. The  indefinite  hydrates  in  solution  may  be  regarded 
as  conforming  to  the  law  of  multiple  proportions,  if  we 
consider  these  hydrates  hi  solution  as  in  a  state  of  dis- 
sociation. 

To  illustrate  and  fix  more  definitely  in  mind  the  views 
of  Mendeleeff  concerning  hydrates,  let  us  look  more 
closely  into  his  work  on  the  hydrates  formed  by  sulphuric 
'acid.1  By  plotting  the  specific  gravities  of  solutions  of 
sulphuric  acid  of  varying  concentrations  against  the  con- 
centrations, curves  were  obtained  with  maxima  hi  them. 
These  maxima  Mendeleeff  interpreted  as  corresponding 
to  the  following  definite  hydrates  —  H2S04,  H2S04.H20, 
H2S04.2H20,  H2S04.25H20,  H2S04.100H20. 

Mendeleeff  points  out  that  the  first  two  of  these  com- 
pounds are  known  in  the  free  state,  the  last  three  are  not 
thus  known.  This  may  be  due  to  a  lack  of  methods  for 
isolating  them,  or  they  may  not  be  capable  of  existing 
in  the  free  condition. 

In  a  similar  manner  Mendeleeff  concluded  that  calcium 
chloride  forms  the  hydrates,  CaCl2.2H20,  CaCl2.4H2O,  and 
CaCl2.6H20.  These  examples  suffice  to  make  clear  the 
idea  which  Mendeleeff  had  in  mind;  certain  compounds, 

1  Ber.  d.  chem.  Gesell,  1,  379  (1886). 


46  THE  NATURE  OF  SOLUTION 

especially  those  that  are  very  hydroscopic  or  have  great 
power  to  combine  with  water,  combine  with  it  when  in 
aqueous  solution,  and  form  a  few  definite  hydrates.  The 
above  examples  will  show  the  composition  of  the  hydrates 
formed  by  a  typical  hygroscopic  acid  and  salt. 

Mendel6eff  furnished  very  little  experimental  evidence 
for  the  existence  of  hydrates  hi  aqueous  solution,  indeed, 
almost  none,  and  saw  no  method  for  determining,  even 
approximately,  the  composition  of  the  hydrates. 

His  so-called  hydrate  theory  was,  therefore,  never 
more  than  a  qualitative  suggestion,  and  furthermore,  has 
been  shown  to  be  erroneous,  at  least  hi  the  case  of  sul- 
phuric acid. 

Testing  Mendeleeffs  Hypothesis.  —  In  the  summer  of 
1893,  the  writer  was  attracted  to  Stockholm  by  the  fame  of 
that  genial  investigator — Svante  Arrhenius.  He  suggested 
that  we  take  up  an  investigation  which,  it  seemed,  would 
test  the  hydrate  theory  of  MendelSeff,  as  far,  at  least, 
as  one  compound  was  concerned.  We  have  already  seen 
how  Mendele'eff  arrived  at  the  conclusion  that  sulphuric 
acid  in  the  presence  of  water  yields  a  few  definite  com- 
pounds with  this  solvent.  This  was  the  conclusion 
which  we  decided  to  test,  and  in  the  following  way.1 

Using  acetic  acid  as  the  solvent,  we  dissolved  hi  it 
known  amounts  of  sulphuric  acid,  and  determined  how 
much  the  freezing-point  of  acetic  acid  was  lowered  by  the 
sulphuric  acid.  We  plotted  the  curve  between  the  varia- 
bles,—  amount  of  sulphuric  acid  and  lowering  of  the 
freezing-point  of  a  definite  amount  of  acetic  acid. 

Similarly,  we  determined  the  lowering  of  the  freezing- 
point  of  a  new  portion  of  acetic  acid  by  varying  amounts 
of  water,  and  plotted  the  curve  between  the  lowering  of 
the  freezing-point  and  the  amount  of  water  present. 

We  then  determined  the  lowering  of  the  freezing- 
point  of  acetic  acid  produced  by  a  known  amount  of 
sulphuric  acid  and  a  known  amount  of  water,  simul- 

1  Zeit.  phys.  Chem.,  13,  419  (1894);  Amer.  Chem.  Journ.,  16,  1  (1894). 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION        47 

taneously.  If  there  is  no  combination  between  the  water 
and  the  sulphuric  acid,  the  lowering  in  the  last  case 
should  be  the  sum  of  the  lowerings  hi  the  first  two  cases. 
If  there  is  combination,  the  lowering  produced  by  the  sul- 
phuric acid  and  water  simultaneously  would  be  less  than 
the  sum  of  the  lowerings  produced  by  them  individually. 

The  latter  was  found  to  be  the  case.  When  sulphuric 
acid  and  water  were  brought  simultaneously  into  the  pres- 
ence of  the  acetic  acid,  the  freezing-point  lowering  was 
less  than  the  sum  of  the  lowerings  when  they  were  added 
independently.  This  showed  that  there  was  some  com- 
bination between  the  sulphuric  acid  and  the  water. 

By  comparing  the  three  sets  of  lowerings  quantitatively, 
it  is  a  very  simple  matter  to  calculate  the  amount  of  water 
hi  combination  with  the  sulphuric  acid.  Knowing  the 
total  amount  of  sulphuric  acid  present  and  the  total 
amount  of  water,  it  is  easy  to  calculate  the  number  of 
molecules  of  water  in  combination  with  one  molecule  of 
sulphuric  acid. 

We  found  that  sulphuric  acid  combines  with  water  form- 
ing the  hydrates  —  H2S04.H20  and  H2S04.2H20.  When  as 
.much  as  thirty-seven  equivalents  of  water  were  present  to 
one  of  sulphuric  acid,  there  was  no  evidence  of  the  forma- 
tion of  hydrates  with  more  water  than  H2S04.2H20,  i.e., 
H6S06.  The  higher  hydrates  which  Mendeleeff  thought  to 
be  probable,  were  not  found  to  be  present  under  the  con- 
ditions under  which  we  worked.  If  they  are  capable  of 
existence,  it  would  seem  that  some  indication  of  them 
should  have  manifested  itself  in  this  work. 

Mendeleeff,1  from  the  specific  gravity  of  solutions  of 
alcohol  in  water,  concluded  that  alcohol  is  capable  of  form- 
ing the  hydrates,  3C2H6O.H20,  C2H60.3H20,  and  C2H60.- 
12H20.  This  was  tested  just  as  the  sulphuric  acid  had 
been.  Acetic  acid  was  used  as  the  solvent,  and  the  lower- 
ing of  its  freezing-point  by  water  and  by  alcohol  separately 
and  when  together,  determined.  By  comparing  the  three 

i  Zeit.  phys.  Chem.,  1,  284  (1887). 


48  THE  NATURE  OF  SOLUTION 

sets  of  results  we  may  decide  whether  the  water  and  the 
alcohol  were  in  a  state  of  combination,  and  if  so,  the 
composition  of  the  hydrates  formed. 

The  results  were  unambiguous,  and  showed  that  there 
is  not  the  slightest  evidence  of  any  hydrate  of  alcohol 
being  formed  under  the  conditions  of  this  experiment. 

The  results  of  this  work  as  a  whole  are  all  against  the 
correctness  of  the  Mendele*eff  theory  of  hydrates.  Indeed, 
the  evidence  which  has  been  obtained  in  this  laboratory 
during  the  past  fifteen  years,  some  of  which  is  discussed 
in  the  closing  chapters  of  this  book,  is  so  decidedly  against 
it  as  a  quantitative  expression  of  the  facts,  that  it  can  no 
longer  be  regarded  as  tenable. 

End  of  the  Qualitative  Period.  —  This  ends  what  may 
be  called  the  qualitative  period  in  the  study  of  solutions. 
Several  investigators  had  supposed  that  there  is  combina- 
tion between  the  solvent  water  and  the  dissolved  sub- 
stance, but  no  one  had  found  any  method  of  determining 
the  composition  of  the  hydrates  formed.  None  of  them 
had  even  furnished  any  very  conclusive  evidence  of  the 
existence  of  these  hydrates.  No  one  up  to  this  tune  had 
supposed  that  combination  between  solvent  and  solute 
is  a  general  phenomenon,  more  or  less  independent,  except 
in  magnitude,  of  the  nature  of  the  solvent  and  of  the 
dissolved  substance. 

As  to  the  relation  between  matter  in  the  dissolved  and 
hi  the  gaseous  state,  what  had  been  pointed  out  was 
purely  qualitative.  As  we  have  seen,  Gay-Lussac  sup- 
posed that  there  was  some  analogy  between  the  two 
states;  Horstmann,  and  Guldberg  and  Waage  made  this 
relation  still  more  probable,  and  Thomsen  laid  consider- 
able stress  upon  it.  Mendeleeff  saw  a  relation  between  the 
solute  in  very  dilute  solutions  and  matter  in  the  gaseous 
state;  but  it  remained  for  another,  Van't  Hoff,  as  we 
shall  see,  to  deduce  this  relation  mathematically  and  thus 
place  it  upon  a  quantitative  basis. 

The   English    chemist,    Williamson,    and   the   German 


EARLIER  VIEWS  OF  THE  NATURE  OF  SOLUTION       49 

physicist,  Clausius,  supposed  that  in  solutions,  acids,  bases, 
and  salts,  at  least,  were  more  or  less  broken  down  into 
their  part  molecules.  Valson  was  able  to  explain  'the 
additive  properties  of  dilute  solutions  of  salts  only  on  the 
same  assumption;  and  Raoult,  especially  from  his  inves- 
tigations of  the  lowering  of  the  freezing-point  and  of  the 
vapor-tension  of  the  solvent  by  the  dissolved  substance, 
was  forced  to  conclude  that  the  additive  property  of  the 
constituents  of  the  dissolved  substances  could  be  explained 
only  on  the  assumption  that  these  constituents  were  prac- 
tically free  from  one  another  in  the  solution.  This  again 
was  purely  qualitative.  It  remained  for  the  Swedish 
physicist,  Arrhenius,  to  prove  the  correctness  of  this 
conclusion,  and  to  furnish  quantitative  methods  for  meas- 
uring the  magnitude  of  this  dissociation.  However,  before 
turning  to  these  quantitative  generalizations,  we  must 
study  a  property  of  solutions  upon  which  these  epoch- 
making  generalizations  rest,  viz.,  their  osmotic  pressure. 


CHAPTER  III 

THE  OSMOTIC  PRESSURE  OF  SOLUTIONS 

Traube's  Method  of  Measuring  Osmotic  Pressure. — 
It  is  not  a  difficult  matter  to  demonstrate  the  existence 
in  solution  of  what  is  known  as  osmotic  pressure.  It  is 
only  necessary  to  close  the  end  of  a  glass  tube  with  parch- 
ment paper,  put  a  solution  of  cane  sugar  in  the  tube,  and 
plunge  the  whole  into  pure  water.  The  solution  cannot 
pass  out  through  the  membrane.  Water  will  pass  in 
and  will  rise  in  the  tube  producing  hydrostatic  pressure. 
While  it  is  such  a  simple  matter  to  demonstrate  the  exis- 
tence of  osmotic  pressure,  it  is  a  very  difficult  matter, 
indeed,  to  measure  it.  Were  we  dependent  on  natural 
membranes  for  measuring  osmotic  pressure  we  would 
know  far  less  about  it  at  present  than  we  do.  For- 
tunately this  is  not  the  case. 

In  1867  Mauritz  Traube  prepared  artificial  mem- 
branes, which  have  the  same  osmotic  properties  as  the 
natural  membranes  referred  to  above.  When  a  sub- 
stance like  copper  ferrocyanide,  formed  by  the  action  of 
potassium  ferrocyanide  on  a  solution  of  a  copper  salt,  is 
deposited  upon  a  resistant  support,  it  manifests  the 
property  of  semipermeability,  as  we  say,  i.e.,  allows  water 
to  pass  through,  but  does  not  allow  the  dissolved  sub- 
stance to  pass.  The  resistant  support  that  is  generally 
employed  is  an  unglazed  porcelain  cup,  the  gelatinous 
precipitate  being  deposited  right  in  the  wall  of  the  cup,  or 
on  its  surface. 

Pfeffer's  Preparation  of  Semipermeable  Membranes. — 
Utilizing  the  discovery  of  M.  Traube,  Pfeffer  was  able  to 
prepare  artificial  semipermeable  membranes  which  would 
not  only  withstand  considerable  pressures,  amounting  to 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  51 

several  atmospheres,  but  which  could  be  used  to  measure 
these  pressures  with  a  fair  degree  of  approximation. 

The  work  of  Pfeffer,1  published  in  1877,  contains  the 
first  quantitative  measurements  of  osmotic  pressure,  worthy 
of  the  name,  that  had  ever  been  made.  Although  more 
recent  work  shows  that  the  results  of  Pfeffer  contain  very 
considerable  errors,  amounting  to  as  much  as  several  per- 
cent, yet,  when  we  consider  that  they  were  the  pioneer 
results  hi  this  field,  the  investigation  which  led  to  them 
must  be  regarded  as  of  very  great  value.  And,  further, 
when  we  see  to  what  the  results  of  Pfeffer  led,  we  are 
almost  inclined  to  regard  this  as  one  of  the  classical 
investigations  of  the  period  in  which  it  was  carried  out. 

A  brief  discussion  of  the  method  employed  by  Pfeffer 
for  making  semipermeable  membranes  can  give  no  idea 
of  the  difficulties  which  he  encountered  and  overcame.  It 
required  several  years  before  he  could  get  even  one 
porcelain  factory  in  Germany  to  prepare  the  unglazed 
porcelain  cups  sufficiently  fine-grained  and  homogeneous 
for  his  purpose.  After  suitable  cups  had  been  obtained, 
the  preparation  of  membranes  which  would  withstand 
-the  pressure  without  leaking  required  years  of  trial.  The 
only  way  in  which  to  obtain  any  adequate  idea  of  what 
Pfeffer  did,  and  what  he  finally  succeeded  in  accom- 
plishing, is  to  consult  his  monograph,  referred  to  above. 

Pfeffer's  Measurements.  —  After  a  suitable  membrane 
had  been  laid  down  hi  the  cup,  the  "cell,"  as  it  was 
called,  was  then  ready  for  use  in  measuring  osmotic  pres- 
sure. The  solution  whose  osmotic  pressure  was  to  be 
measured  was  placed  hi  the  cup,  which  was  then  closed 
with  a  stopper  having  a  manometer  attached.  The  cup 
was  then  placed  in  pure  water,  and  after  equilibrium  had 
been  established  the  pressure  read  on  the  manometer. 

Pfeffer  used  several  gelatinous  precipitates  as  semi- 
permeable  membranes.  Thus,  he  deposited  in  the  walls 
of  the  cups  Berlin  blue,  calcium  phosphate,  etc.,  but  found 

1  Osmotische  Untersuchungen,  Leipzig,  1877. 


52  THE  NATURE  OF  SOLUTION 

that  these  substances  would  not  withstand  the  pressure 
as  well  as  copper  ferrocyanide.  Indeed,  they  would  break 
before  the  maximum  pressure  was  reached.  All  of  Pfeffer's 
best  measurements  were  therefore  made  with  copper 
ferrocyanide  as  the  semipermeable  membrane. 

Pfeffer's  Results. —  Pfeffer  determined  the  osmotic  pres- 
sures of  five  solutions  of  cane  sugar.  His  results  are  given 
in  the  following  table  taken  from  "Osmotische  Untersuch- 
ungen,"  p.  81. 

Concentrations  in  Osmotic  pressures 

percent  by  weight  in  centimeters  of  mercury 

1  53.5  cm. 

2  101.6  " 
2.74  151.8  " 
4  208.2  " 
6  307.5  " 

Gum  arabic  gave  much  smaller  osmotic  pressures. 

Concentrations  in  Osmotic  pressures 

percent  by  weight  in  centimeters  of  mercury 

1  7.1  cm. 

6  27.5    " 

18  120.0    " 

Pfeffer  also  measured  the  osmotic  pressures  of  a  few 
solutions  of  one  electrolyte — potassium  nitrate.  He  started 
with  solutions  having  the  concentrations,  1,  2,  and  4  weight 
percent  respectively. 

It  was  impossible  for  him  to  find  a  membrane  through 
which  this  substance  would  not  pass  to  some  extent. 
There  was  a  constant  diminution  in  the  strength  of  the 
solution,  due  to  the  salt  dialyzing  through  the  membrane. 
He  determined  the  concentrations  of  the  solutions  just 
after  reading  their  osmotic  pressures,  and  obtained  the 
following  results 1  —  the  above  concentrations  having 
changed  to  0.80.  1.43,  and  3.3  percent,  respectively. 

Concentrations  in  Osmotic  pressures 

weight  percent  in  centimeters  of  mercury 
0.80  130.4  cm. 

0.86  147.5    " 

1.43  218.5    " 

3.30  436.8    " 

1  Osmotische  Untersuchungen,  p.  82. 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  53 

These  results  were  all  obtained  at  one  temperature. 
Pfeffer  measured  also  the  osmotic  pressures  of  certain 
solutions  of  different  substances  at  several  temperatures. 
These  results  are  important  from  the  standpoint  of  the 
nature  of  solution,  as  we  shall  see  later.  A  few  of  his 
results  are  given  below. 

The  osmotic  pressures  of  a  one  percent  solution  of  cane 
sugar  were  measured  at  several  temperatures  between  6.8° 
and  36°,  and  the  following  results  obtained.1 

Temperatures  Osmotic  pressures 

0  /  14.2°  51.0  cm. 

a\32.0°  54.4    " 

f   6.8°  50.5    " 

b    13.7°  52.5    " 

[22.0°  54.8    " 

f!5.5°  52.0    " 

c\36.0°  56.7    " 

The  apparent  irregularities  are  due  to  the  fact  that 
different  cells  were  used  hi  the  several  measurements. 
These  are  marked  (a),  (b),  and  (c).  Only  those  results 
obtained  in  the  same  cell  are  comparable  with  one  another. 

The  following  results  were  obtained  with  a  0.6  percent 
solution  of  sodium  tartrate. 

Temperatures  Osmotic  pressures 

12.4°  91.6  cm. 

37.3°  98.3    " 

A  saturated  solution  of  acid  potassium  tartrate  gave, 
at  two  temperatures,  the  following  results: 

Temperatures  Osmotic  pressures 

13.0°  68.3  cm. 

29.2°  115.8    " 

These  are  only  some  of  the  results  obtained  by  Wil- 
helm  Pfeffer. 

Where  Pfeffer  Left  the  Subject  of  Osmotic  Pressure. — 
Pfeffer  was  a  botanist,  and  took  up  the  problem  of  osmotic 

1  Ibid.,  p.  85. 


54  THE  NATURE  OF  SOLUTION 

pressure  from  a  botanical  standpoint.  He  wanted  to 
know  something  of  the  magnitude  of  the  force  which  gives 
us  osmotic  pressure.  Is  it  a  very  small  force,  or  is  it  one 
of  considerable  magnitude?  His  own  results  answered 
this  question.  A  one  percent  solution  of  cane  sugar 
showed  an  osmotic  pressure  of  53  centimeters  of  mercury  — 
nearly  two-thirds  of  an  atmosphere.  A  one  percent  solu- 
tion contains  ten  grams  in  a  liter.  The  molecular  weight 
of  cane  sugar  is  342.  Ten  grams  in  a  liter  is,  therefore, 
only  one  thirty-fourth  normal.  A  normal  solution  would, 
from  the  work  of  Pfeffer,  have  an  osmotic  pressure  of 
about  twenty-two  atmospheres,  which  is  a  force  of  very 
appreciable  order  of  magnitude. 

Osmotic  pressure  undoubtedly  plays  a  prominent  role 
in  plant  life.  How  does  the  water  rise  from  the  bottom 
to  the  top  of  a  tree?  In  the  case  of  the  large  trees  in 
the  West,  this  rise  is  often  as  much  as  three  hundred 
feet;  and  the  amount  of  water  raised  to  this  height  and 
breathed  out  through  the  leaves  is  enormous.  While  this 
problem  cannot  even  yet  be  considered  as  completely 
solved,  it  is  quite  certain  that  osmotic  pressure  has  much 
to  do  with  it. 

The  last  half  of  Pfeffer's  important  monograph1  is 
devoted  to  the  bearing  of  osmotic  pressure  on  plant 
physiology.  It  would  lead  us  beyond  the  scope  of  this 
work  to  discuss  this  part  of  Pfeffer's  work.  Having  solved 
the  problem  which  he  undertook  to  study,  viz.,  the  order 
of  magnitude  of  osmotic  pressure,  Pfeffer  left  it  at  this 
stage.  He  did,  however,  point  out  the  relation  between 
the  osmotic  pressures  of  solutions  and  their  concentrations. 
The  osmotic  pressures  divided  by  the  concentrations  gave 
a  constant.  He  did  not  seem  to  see  anything  very  sig- 
nificant in  this  relation.  Osmotic  pressure  is  simply 
another  of  those  properties  which  is  a  linear  function  of, 
or  proportional  to,  the  concentration  of  the  solution.  This 
amounts  to  saying  that  two  molecules  of  a  dissolved 

1  Osmotische  Untersuchungen,  Leipzig,  1877. 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  55 

substance  exert  twice  the  osmotic  pressure  of  one  mole- 
cule, which  is  not  at  all  surprising. 

To  a  man  with  a  different  type  of  mind,  as  we  shall  see, 
this  relation  meant  far  more  than  simple  proportionality. 
To  Pfeffer,  it  did  not. 

Other  Measurements  of  Osmotic  Pressure. —  There- 
suits  obtained  by  Pfeffer,  while  of  fundamental  importance, 
were  only  close  approximations.  Further,  membranes 
prepared  by  the  Pfeffer  method,  i.e.,  by  allowing  the  two 
solutions  which,  when  they  came  together,  would  form 
the  gelatinous  precipitate,  to  soak  or  diffuse  into  the  walls 
of  the  cell,  the  one  from  the  inside  the  other  from  the 
the  outside,  were  not  capable  of  withstanding  very  great 
pressures.  The  most  concentrated  solution  of  cane  sugar 
whose  osmotic  pressure  Pfeffer  was  able  to  measure,  was 
a  six  percent,  solution.  This  was  only  between  one-fifth 
and  one-sixth  normal,  and  had  an  osmotic  pressure  of  only 
about  four  atmospheres. 

;  These  results  were,  therefore,  not  satisfactory.  They 
were  not  sufficiently  accurate  for  dilute  solutions,  and 
gave  no  direct  clue  to  the  osmotic  pressures  of  even 
moderately  concentrated  solutions.  Both  of  these  prob- 
lems have  now  been  solved.  We  have  far  more  accurate 
measurements  of  the  osmotic  pressures  of  dilute  solu- 
tions, and  approximate  measurements  of  the  osmotic 
pressures  of  concentrated  solutions.  We  shall  now  con- 
sider briefly  this  more  recent  work. 

Work  of  Berkeley  and  Hartley  on  the  Osmotic  Pres- 
sures of  Concentrated  Solutions.  —  The  Earl  of  Berkeley 
and  Hartley1  have  devised  a  method  for  measuring 
roughly  the  osmotic  pressures  of  solutions  up  to  one 
hundred  and  thirty  atmospheres.  Their  method  does  not 
measure  directly  the  osmotic  pressure  of  the  solution,  but 
consists  in  bringing  to  bear  on  the  solution  a  pressure 
which  is  just  equal  and  opposite  to  its  osmotic  pressure, 
and  then  measuring  the  magnitude  of  this  pressure. 

1  Trans.  Roy.  Soc.,  A.,  206,  481  (1906). 


56  THE  NATURE  OF  SOLUTION 

The  semipermeable  membrane  of  copper  ferrocyanide 
is  deposited  on  the  walls  of  an  unglazed  porcelain  tube.  The 
porcelain  tube  is  surrounded  by  a  tube  of  gun  metal  to  which 
a  pressure  apparatus  is  attached.  Water  is  introduced  into 
the  porcelain  tube  containing  on  its  walls  the  gelatinous 
precipitate  serving  as  a  semipermeable  membrane.  The 
solution  whose  osmotic  pressure  is  to  be  measured  is  intro- 
duced into  the  gun-metal  tube,  and  therefore  comes  hi 
contact  with  the  porcelain  tube  and  the  semipermeable 
membrane  on  the  outside. 

Water  would  pass  from  the  porcelain  tube  through  the 
semipermeable  membrane  into  the  solution,  but  a  counter- 
pressure  is  brought  to  bear  mechanically  upon  the  solution, 
and  this  is  increased  in  magnitude  until  water  appears  to 
cease  to  flow  through  the  walls  of  the  porcelain  tube  into 
the  solution.  This  is  determined  by  having  a  capillary 
glass  tube  attached  to  the  inner  side  of  the  porcelain  tube 
and  observing  whether  the  height  of  the  water  in  this  tube 
just  remains  stationary.  If  the  level  of  the  water  in  the 
capillary  falls,  it  means  that  water  is  passing  through  the 
membrane  into  the  solution,  and  more  mechanical  pressure 
must  be  brought  to  bear  on  the  solution  to  prevent  the  water 
from  entering.  If  the  water  rises  in  the  capillary,  it  means 
that  water  is  passing  through  the  membrane  from  the  so- 
lution, showing  that  the  mechanical  pressure  on  the  solu- 
tion is  greater  than  its  osmotic  pressure.  The  pressure  on 
the  solution  must  be  so  adjusted  that  the  level  of  the  water 
in  the  capillary  neither  rises  nor  falls.  The  mechanical 
pressure  is  then  just  equal  to  the  osmotic  pressure  of  the 
solution.  It  is  then  only  necessary  to  measure  the  pressure 
brought  to  bear  on  the  solution,  and  we  know  its  osmotic 
pressure. 

Results  Obtained.  —  A  few  of  their  results  are  given 
below  —  cane  sugar:  (molecular  weight  342).  The  con- 
centrations are  given  hi  grams  per  liter.  The  temperature 
was  0°. 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  57 

Grams  cane  sugar  Osmotic  pressures 

in  a  liter  of  solution  in  atmospheres 

180.1  13.95 

300.2  26.77 

420.3  43.97 

540.4  67.51 

660.5  100.78 

750.6  133.74 

The  following  results  were  obtained  with  dextrose.1 

Grams  dextrose  in  a  Osmotic  pressures 

liter  of  solution  in  atmospheres 

99.8  13.21 

199.5  29.17 
319.2  53.19 

448.6  87.87 
548.6  121.18 

The  results  with  mannite  are: 

Grams  mannite  in  a  Osmotic  pressures 

liter  of  solution  in  atmospheres 

100  13.1 

110  14.6 

125  16.7 

Five  hundred  grams  of  galactose  hi  a  liter  gave  an  osmotic 
pressure  of  95.8  atmospheres. 

These  results  must  be  regarded  as  only  approximations, 
the  necessary  error  involved  being  large.  They  are,  how- 
ever, of  the  right  order  of  magnitude  and  are  doubtless  very 
nearly  the  osmotic  pressures  of  these  concentrated  solu- 
tions. We  are  impressed  by  their  large  magnitude.  Take 
the  most  concentrated  solution  of  cane  sugar,  which  contains 
750.6  grams  per  liter.  The  molecular  weight  of  cane  sugar 
being  342,  this  solution  is  about  2.2  normal.  —  The  osmotic 
pressure  of  a  normal  solution  of  cane  sugar  is,  as  we  shall 
see,  between  twenty-five  and  thirty  atmospheres,  depending 
on  the  temperature.  The  osmotic  pressure  of  a  two  normal 
solution  of  cane  sugar  is  therefore  more  than  three  times 
that  of  a  normal.  The  meaning  of  this  rapid  increase  hi  the 
osmotic  pressure  with  concentration  will  become  apparent 
when  we  come  to  consider  the  solvate  theory  of  solution. 

1  Trans.  Roy.  Soc.,  A.,  206,  503  (1906). 


58  THE  NATURE  OF  SOLUTION 

Pfeffer's  Membranes  Would  Not  Withstand  Great  Pres- 
sures. —  Pfeffer,  as  will  be  remembered,  filled  the  cells  with 
a  solution  of  potassium  ferrocyanide,  and  then  plunged  them 
into  a  solution  of  copper  sulphate.  The  two  solutions  dif- 
fused into  the  walls  of  the  cup,  the  one  from  the  inside,  the 
other  from  the  outside.  Where  the  copper  from  the  copper 
sulphate  came  in  contact  with  the  ferrocyanogen  group 
from  the  ferrocyanide,  a  precipitate  of  copper  ferrocyanide 
was  formed.  By  allowing  the  solution  to  stand  in  the  cup 
for  the  proper  length  of  tune  before  plunging  the  cup  into 
the  copper  sulphate,  the  precipitate  would  be  formed  on  the 
inner  surface  of  the  cup  or  anywhere  in  its  wall. 

Pfeffer,  as  will  be  seen,  had  to  rely  entirely  upon  dif- 
fusion to  form  his  membranes,  and,  as  would  be  expected,  the 
membranes  would  be  only  fairly  resistant  to  pressure.  They 
would  not  withstand  great  pressure;  indeed,  only  pressures 
of  a  few  atmospheres. 

Electrical  Method  of  Preparing  Semi-permeable  Mem- 
branes.—  A  method  of  making  semipermeable  membranes 
was  discovered  in  this  laboratory,  by  H.  N.  Morse,  which 
gives  results  far  more  satisfactory  than  that  employed 
by  Pfeffer.  Instead  of  allowing  the  membrane  formers  to 
soak  into  the  walls  of  the  cell,  they  are  driven  hi  electrically. 
The  cell  is  filled  with  a  solution  of  potassium  ferrocyanide, 
and  plunged  into  a  solution  of  copper  sulphate.  One  elec- 
trode is  placed  in  the  cell,  and  the  other  around  the  cell  hi 
the  solution  of  copper  sulphate.  The  current  is  passed  from 
the  solution  of  copper  sulphate  to  that  of  potassium  ferro- 
cyanide. The  copper  ion,  moving  with  the  current,  passes 
into  the  walls  of  the  cell  from  the  outside;  the  ferro- 
cyanogen group,  moving  against  the  current,  passes  into 
the  walls  of  the  cell  from  the  inside.  Where  the  two  meet 
they  form  copper  ferrocyanide. 

The  advantage  of  this  method  over  the  method  of 
Pfeffer  is,  that  by  driving  the  membrane  formers  into  the 
walls  electrically,  a  much  more  resistant  membrane  can  be 
produced  than  by  simply  allowing  them  to  soak  into  the 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  59 

walls.  The  current  is  passed  until  the  membrane  becomes 
so  compact  that  the  resistance  may  rise  to  one  million 
ohms.  Membranes  formed  by  this  method  have  been  used 
to  measure  osmotic  pressures  of  thirty  and  recently  as  great 
as  two  hundred  and  fifty  atmospheres,  with  a  high  degree 
of  accuracy. 

It  was  the  discovery  of  this  method  of  preparing  the 
membranes  that  led  Morse  to  undertake  anew  the  measure- 
ment of  osmotic  pressure;  the  result  being  that  he  and  his 
co-workers  have  made  measurements  of  osmotic  pressure 
incomparably  the  best. 

Measurements  of  Osmotic  Pressure  Made  by  Morse, 
Frazer,  Holland  and  Co-workers.  —  Provided  with  the 
above  described  method  of  making  sernipermeable  mem- 
branes, Morse  and  his  co-workers  began  about  fifteen  years 
ago  a  series  of  measurements  of  the  osmotic  pressures  of 
solutions,  which,  for  the  difficulties  encountered  and  over- 
come, and  for  the  accuracy  of  the  results  obtained,  have 
become  a  classic. 

For  details  in  connection  with  this  work  reference  must 
be  had  to  the  original  publications.1  They  have  measured 
thus  far  the  osmotic  pressures  of  solutions  of  cane  sugar 
varying  hi  concentration  from  one-tenth  to  normal,  and  at 
the  Mowing  temperatures,  —  0°,  5°,  10°,  15°,  20°,  25°, 
30°,  40°,  50°,  60°,  70°,  and  80°.  The  solutions  were  all 
prepared  on  a  weight  normal  basis,  i.e.,  so  many  grams  of 
sugar  in  one  thousand  grams  of  solvent. 

The  mean  osmotic  pressures  of  solutions  of  cane  sugar 
from  one-tenth  to  normal,  and  from  0°  to  80°,  are  recorded 
hi  the  following  table. 

1  Amer.  Chem.  Journ.-,  during  the  past  fifteen  years.  Carnegie  Institu- 
tion of  Washington,  Publication  No.  198. 


60 


THE  NATURE  OF  SOLUTION 


RESULTS  WITH  CANE  SUGAR 

OSMOTIC  PRESSURES  IN  ATMOSPHERES 
FOR  WEIGHT-NORMAL  CONCENTRATIONS 


TEMP. 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

0.7 

0.8 

0.9 

1.0 

0° 

(2.462) 

4.723 

7.085 

9.443 

11.895 

14.381 

16.886 

19.476 

22.118 

24.826 

5° 

2.452 

4.819 

7.198 

9.608 

12.100 

14.605 

17.206 

19.822 

22.477 

25.280 

10° 

2.498 

4.893 

7.335 

9.790 

12.297 

14.855 

17.503 

20.161 

22.884 

25.693 

15° 

2.540 

4.985 

7.476 

9.949 

12.549 

15.144 

17.815 

20.535 

23.305 

26.189 

20°^ 

2.590 

5.064 

7.605 

10.137 

12.748 

15.388 

18.128 

20.905 

23.717 

26.638 

25° 

2.634 

5.148 

7.729 

10.296 

12.943 

15.625 

18.435 

21.254 

24.126 

27.053 

30° 

2.474 

5.044 

7.647 

10.295 

12.978 

15.713 

18.499 

21.375 

24.226 

27.223 

40° 

2.560 

5.163 

7.834 

10.599 

13.355 

16.146 

18.932 

21.806 

24.735 

27.701 

50° 

2.637 

5.279 

7.974 

10.724 

13.504 

16.314 

19.202 

22.116 

25.123 

28.209 

60° 

2.717 

5.438 

8.140 

10.866 

13.666 

16.535 

19.404 

22.327 

25.266 

28.367 

70° 

13.991 

16.820!  19.568 

22.567 

25.562 

28.624 

80° 

23.062 

25.919 

28.818 

Ratios  Between  Gas  Pressure  and  Osmotic  Pressure.  — 
As  we  shall  see  in  the  next  chapter,  the  importance  of  os- 
motic pressure  as  bearing  on  the  nature  of  solution  is  prima- 
rily hi  connection  with  the  relation  between  these  pressures 
in  solutions  and  the  gas-pressure  of  gases  having  the  same 
number  of  gaseous  parts  hi  a  given  volume  that  there  are 
dissolved  parts  in  the  same  volume  of  the  solution.  With- 
out going  more  into  detail  here  on  this  point,  it  should  be 
stated  that  Morse  and  his  co-workers  have  compared  the 
osmotic  pressures  of  solutions  of  cane  sugar  with  the  gas- 
pressures  of  gases;  "the  volume  of  the  gas  being  that  of  the 
solvent  in  the  pure  state ,  and  not  that  of  the  solution."2 

The  ratios  of  osmotic  to  gas-pressure  are  given  by 
Morse  hi  the  following  table:3 

1  These  results  are  taken  from  Carnegie  Institution  of  Washington,  Pub- 
lication No.  198,  pp.  184  and  186  (1914). 

•  Ibid.,  p.  183.  3  nrid.,  p.  186. 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS 


61 


RATIO  OF  OSMOTIC  TO  GAS-PRESSURE 


TEMP. 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

0.7 

0.8 

0.9 

1.0 

0° 

1.106 

1.061 

1.061 

1.060 

1.068 

1.076 

1.083 

1.093 

.104 

1.115 

5° 

1.082 

1.063 

1.058 

1.059 

1.067 

1.074 

1.084 

1.093 

.102 

1.115 

10° 

1.082 

1.060 

1.059 

1.060 

1.066 

1.073 

1.083 

1.092 

.102 

1.113 

15° 

1.082 

1.061 

1.061 

1.059 

1.068 

.073 

1.083 

1.093 

.102 

.115 

20° 

1.084 

1.062 

1.060 

.060 

1.067 

.073 

1.084 

1.093 

.103 

.115 

25° 

1.084 

1.059 

1.060 

.059 

1.065 

.071 

1.083 

.093 

.102 

.113 

30° 

1.000 

1.020 

1.031 

.040 

1.050 

.060 

1.069 

.081 

.089 

.101 

40° 

1.003 

1.011 

1.024 

.038 

1.046 

.054 

1.059 

.067 

.076 

.085 

50° 

1.000 

1.00? 

1.009 

.017 

1.025 

.032 

1.041 

.049 

.059 

.071 

60° 

1.000 

1.001 

0.999 

1.000 

1.006 

1.014 

1.020 

.027 

1.033 

1.044 

70° 

1.000 

1.002 

0.999 

.008 

1.015 

1.023 

80° 

1.001 

1.000 

1.000 

"The1  ratio  of  osmotic  to  gas  pressure,  between  0°  and 
25°,  is  very  nearly  constant  for  each  (Concentration  of  solution. 
Omitting  the  ratio  for  the  0.1  weight-normal  solution  at 
0°,  the  mean  ratios  for  the  various  concentrations  are  given 
below." 

Concentration  Mean  ratio 

weight  normal 

0.1 

0.2 

0.3 


0.4 
0.5 
0.6 
0.7 
O.S 
0.9 
1.0 


.083 
.061 
.060 
.060 
.067 
.074 
.083 
.093 
.103 
.114 


"The2  average  deviation  from  these  mean  ratios  is  0.15 
percent.,  while  the  largest  single  deviation  —  that  of  the 
0.6  normal  solution  at  25°,  is  0.3  percent.  It  is  obvious 
from  the  relations  pointed  out  above  that  between  0°  and 
25°  the  osmotic  pressure  of  cane-sugar  solutions,  ranging 
in  concentration  from  0.1  to  1.0  weight-normal,  obeys  the 
law  of  Gay-Lussac  for  gases.  In  other  words,  within  the 
limits  designated  the  temperature  coefficients  of  gas  and  of 
osmotic  pressures  are  identical." 

"The3  ratios  of  osmotic  to  gas  pressure  between  0°  and 


1  Ibid.,  p.  183. 


2  Ibid.,  p.  186. 


3  Ibid.,  p.  186. 


62  THE  NATURE  OF  SOLUTION 

25°,  though  constant  for  each  concentration,  are  all  greater 
than  unity.  The  excess  varies  from  6  per  cent  in  the  0.2, 
0.3,  and  0.4  normal  solutions,  on  the  one  side,  to  8.3  per- 
cent in  the  0.1  normal  solution;  and  on  the  other,  to  11.4 
percent  in  the  normal  solution." 

"  Having  found  that  the  law  of  Gay-Lussac  does  hold 
for  the  osmotic  pressures  of  cane  sugar  solutions  between 
0°  and  25°,  one  is  inclined  to  believe  that  they  should  also 
conform  to  the  law  of  Boyle,  and  to  seek  for  some  rational 
explanation  of  the  facts:  1st,  that  the  ratios  in  question 
are  excessive,  i.e.,  above  unity;  and  2nd,  that  they  are  not 
proportional  to  the  supposed  concentration  of  the  solutions. 
The  most  obvious  general  explanation  (if  one  attempts  to 
reconcile  the  pressures  between  0°  and  25°  to  the  view  that 
the  law  of  Boyle,  as  well  as  that  of  Gay-Lussac,  does  hold) 
is  hydration  of  the  solute,  which  may  be  presumed  to  have  the 
effect  of  concentrating  the  solutions.  But  if  one  attempts 
to  work  out  the  precise  degree  of  hydration  which  would 
account  for  the  variations  of  ratio  from  concentration  to 
concentration,  he  is  quickly  entangled  hi  certain  hazard- 
ous assumptions  respecting  the  relations  of  solvent  to  solute 
and  the  effect  of  these  upon  the  osmotic  pressure.  In 
the  writer's  opinion,  judgment  as  to  the  applicability  of 
Boyle's  law  to  the  osmotic  pressure  of  cane-sugar  solu- 
tions at  temperatures  below  25°  should  be  suspended  until 
much  more  is  known  about  the  osmotic  pressures  of  the 
aqueous  solutions  of  other  substances." 

Summary  of  Relations.  —  Morse  summarizes  these  rela- 
tions as  follows:1  " Between  0°  and  25°  the  ratios  of  os- 
motic to  gas  pressure  are  all  greater  than  unity,  but  constant 
for  each  concentration.  At  some  temperature  between 
25°  and  30°  these  ratios  begin  to  decline,  but  relatively 
more  rapidly  in  the  dilute  than  in  the  concentrated  solu- 
tions. At  some  temperature  (30°  for  0.1;  50°  for  0.2;  60° 
for  0.3  and  0.4;  70°  for  0.5,  0.6,  and  0.7;  and  80°  for  0.8, 
0.9,  and  1.0)  the  ratio  becomes  unity  for  every  concentration." 

1  Carnegie  Institution  of  Washington,  Publication  No.  198,  p.  187. 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  63 

"The  decrease  in  the  ratios  of  osmotic  to  gas  pressures 
at  temperatures  above  25°,  suggests  an  increasing  dilution  of 
the  solutions  through  the  dissociation  of  unstable  hydrates; 
and  it  serves  to  strengthen  the  impression  that  the  exces- 
sive but  constant  ratios  below  25°  are  due  to  the  presence  of 
stable  hydrates."  ...  "It  is  now  important  to  ascertain 
whether  the  ratios,  having  once  declined  to  unity,  maintain 
that  value  at  all  higher  temperatures;  hence  the  work  of 
measuring  the  osmotic  pressure  of  cane  sugar  at  70°  and  80°, 
and  at  still  higher  temperatures,  will  be  resumed  as  soon  as 
the  new  cells,  previously  referred  to,  have  been  sufficiently 
developed  for  use  at  those  temperatures." 

It  should  be  mentioned  in  connection  with  this  investiga- 
tion, that  a  cell  containing  a  0.5  weight-normal  solution 
was  allowed  to  stand  for  sixty  days  after  the  maximum 
osmotic  pressure  had  manifested  itself,  and  at  the  end  of 
this  period  showed  no  weakening  of  the  membrane. 

Results  with  Glucose.  —  Morse  and  his  co-workers 
measured  the  osmotic  pressures  of  a  series  of  solutions  of 
glucose  of  varying  concentration.  They  worked  between 
0.1  and  1.0  normal,  and  at  30°  and  40°  and  50°.  It  would 
lead  us  too  far  to  give  the  results1  hi  detail  here.  We  shall 
simply  quote  the  conclusions  reached.2  "The  foregoing 
measurements  of  the  osmotic  pressures  of  glucose  indicate 
that,  between  30°  and  50°,  the  aqueous  solutions  of  this 
substance  obey  the  gas  laws,  since  —  if  we  employ  the 
weight  or  solvent  normal  system  hi  making  the  solutions, 
and  refer  the  theoretical  gas  pressure  of  the  solute  to  the 
volume  of  the  pure  solvent  —  the  ratio  of  observed  osmotic 
to  calculated  gas  pressure  is,  in  all  cases,  approximately 
unity.  Stated  hi  another  way,  the  equation  of  Van't 
Hoff  for  very  dilute  solutions,  PV  =  RT,  applies  to  concen- 
trated solutions  of  glucose  between  30°  and  50°,  provided 
we  allow  the  V  to  signify  the  volume  of  the  pure  solvent 
instead  of  the  volume  of  the  solution." 

1  Ibid.,  p.  196.  >  Ibid.,  p.  207. 


64  THE  NATURE  OF  SOLUTION 

The  meaning  of  all  this  will  become  the  more  apparent 
when  the  relations  pointed  out  in  the  next  chapter  are  con- 
sidered. 

Results  with  Mannite.  —  From  the  results  with  mannite 
the  following  conclusion  was  drawn.  "It  will  be  seen  that 
all  the  ratios  approach  unity,  showing  that  within  the  limits 
thus  far  investigated  the  aqueous  solutions  of  mannite 
obey  the  laws  of  Gay-Lussac  and  of  Boyle." 

Osmotic  Pressure  of  an  Electrolyte.  —  Morse  and  Frazer 
have  measured  the  osmotic  pressures  of  a  few  solutions  of 
one  electrolyte — lithium  chloride,  between  the  concentra- 
tions 0.1  and  0.6  normal;  the  temperature  being  30°. 

The  enormous  difficulties  which  were  encountered  in 
measuring  the  osmotic  pressures  of  solutions  of  non-elec- 
trolytes, were  greatly  augmented  when  the  attempt  was 
made  to  measure  the  osmotic  pressure  of  any  solution  of  an 
electrolyte.  The  reason  for  this  will  become  clear,  when  we 
consider  the  action  of  an  electrolyte  on  a  colloid  like  the 
copper  ferrocyanide  used  for  the  semipermeable  membrane. 
When  we  come  to  study  colloidal  solutions,  we  shall  see  that 
when  an  electrolyte  (acid,  base,  and  especially  salt)  is  brought 
in  contact  with  a  colloid,  the  colloid  crystallizes  and  loses 
its  colloidal  properties.  It  was  found  that  when  a  cell 
containing  a  good,  semipermeable  membrane  was  filled 
with  a  solution  of  an  electrolyte,  the  membrane  rapidly 
degenerated  and  soon  began  to  leak.  This  was  probably 
due  to  the  conversion  of  the  colloidal  membrane  into  the 
crystalline  condition.  Sodium  and  potassium  chlorides  were 
so  vigorous  hi  their  action  on  the  membrane,  that  the  os- 
motic pressure  of  even  dilute  solutions  of  these  substances 
could  not  be  measured.  Lithium  chloride  in  fairly  concen- 
trated solution  also  caused  the  membrane  to  degenerate, 
but  in  dilute  solution  its  action  was  so  slow  that  the  osmotic 
pressures  of  several  solutions  of  lithium  chloride  were  meas- 
ured with  a  fair  degree  of  accuracy. 

It  is  a  matter  of  interest  to  know  that  membranes  which 
have  deteriorated  under  the  action  of  electrolytes  can  be 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS 


65 


restored,  in  part  at  least,  by  standing  a  long  time  in  contact 
with  pure  water.  Some  two  years  are  required  to  effect 
anything  like  complete  restoration.  The  probable  action 
of  the  water  is  to  restore  the  colloidal  condition  of  the  mem- 
brane; and  this  raises  the  question  whether  water  may  not 
have  the  general  tendency  to  transform  crystalloids  into 
colloids.  The  small  number  of  charged  parts  contained 
hi  it  would  make  this  possible,  charged  parts,  as  we  shall 
see,  tending  to  convert  colloids  into  crystalloids. 

Results  with  Lithium  Chloride.  —  The  following  results 
were  obtained  for  the  osmotic  pressures  of  solutions  of 
lithium  chloride,  ranging  hi  concentration  from  0.1  to  0.6 
normal,  the  temperature  being  30°. 1 

CONCENTRATIONS 


0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

Observed  pressures  

4.325 
4311 

8.946 
9005 

13.809 
13  626 

18.755 
18789 

24.162 

29.535 

Mean  pressures 

4.317 

8.976 

13.768 

18.772 

24.162 

29.535 

Calculated  gas  pressures  .... 
Ratio  of  osmotic  to  gas  pres- 
sure 

2.472 
1  746 

4.943 
1  816 

7.415 
1.857 

9.886 
1.899 

12.358 
1.955 

14.830 
1.992 

The  very  large  osmotic  pressures  are  due  to  the  substance 
being  an  electrolyte,  i.e.,  broken  down  in  solution,  as  we  shall 
see,  into  its  part  molecules,  and  each  part  exerting  the  same 
pressure  as  a  whole  molecule. 

That  the  ratio  of  osmotic  to  gas  pressure  increases  with 
the  concentration  of  the  solution,  is  due  to  the  fact  that 
lithium  chloride  is  one  of  those  substances  which  hi  solu- 
tion combines  with  a  large  amount  of  water,  as  was  shown 
in  this  laboratory  a  number  of  years  ago.2  The  more  con- 
centrated the  solution,  the  larger  the  total  amount  of  water 
combined  with  the  dissolved  substance,  and  consequently 
the  smaller  the  percentage  of  the  total  water  present 

1  Taken  from  Carnegie  Institution  of  Washington,  Publication  No.  198, 
p.  217  (1914). 

2  Carnegie  Institution  of  Washington,  Publication  No.  60,  p.  31  (1907). 


66  THE  NATURE  OF  SOLUTION 

acting  as  solvent.  The  solution  is,  therefore,  much  more 
concentrated  than  we  would  suppose,  and  has  a  greater 
osmotic  pressure  and  therefore  a  higher  ratio  between 
osmotic  pressure  and  gas  pressure. 

Durability  of  the  Morse  Cells.  —  The  discussion  of  this 
work  by  Morse  and  his  co-operators  cannot  be  concluded 
better  than  by  quoting  the  following  paragraph,  which  gives 
an  idea  of  the  durability  of  the  cells  which  they  used. 

"  Particular1  attention  is  called  to  experiment  2  with  the 
0.4  weight-normal  solution.  This  is  an  endurance  test  of 
the  membrane  of  an  unusually  thorough  character.  The 
cell  at  the  tune  of  setting  it  up,  had  a  resistance  of  1,100,000 
ohms  and  it  remained  in  the  bath  145  days.  Starting  with 
an  initial  pressure  of  15  atmospheres,  it  reached  an  approxi- 
mate equilibrium  hi  10  days.  The  osmotic  pressure 
which  the  cell  sustained  during  the  following  124  days  is 
given  in  5  columns,  each  of  25  daily  records.  The  mean 
osmotic  pressure  for  the  first  period  was  18,827;  for  the 
second  18,894;  for  the  third  18,799;  for  the  fourth  18,636; 
and  for  the  fifth  18,405.  It  is  believed  that  a  mean  of  the 
records  for  the  first  100  days  fairly  represents  the  osmotic 
pressure  of  the  solution.  But  during  the  fifth  period,  i.e., 
from  the  101st  to  the  125th  day  of  the  record,  there  was  a 
decline  in  pressure  from  18,609  to  18,140  atmospheres,  which 
can  only  signify  that  the  membrane  had  at  last  begun  to 
weaken." 

Greater  durability  than  this  hi  a  semipermeable  mem- 
brane could  neither  be  expected  nor  desired. 

A  Cell  for  the  Measurement  of  High  Osmotic  Pressure. 
—  Frazer,  with  the  co-operation  of  Myrick,  has  devised  a 
cell  for  measuring  the  high  osmotic  pressures  shown  by  con- 
centrated solutions.  This  would  allow  the  laws  of  Boyle 
and  Gay-Lussac  for  osmotic  pressure  to  be  tested  over  a 
wide  range  of  concentration.  A  cell  has  been  constructed 
to  withstand  satisfactorily  pressures  as  high  as  250  atmos- 
pheres. The  position  of  the  membrane  is  changed  in  the 

1  Carnegie  Institution  of  Washington,  Publication  No.  198,  p.  218. 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  67 

cell  earlier  described  from  near  the  interior  surface  of  the 
cell  to  near  the  exterior.  This  same  principle  has  already 
been  successfully  used  by  Berkeley  and  Hartley,  and  is 
necessary  in  order  to  have  greater  strength  hi  the  cell  wall. 
It  also  increases  the  membrane  surface,  which  insures  more 
rapid  adjustment  of  the  equilibrium,  diminishing  the  tune 
for  making  the  measurements.  In  constructing  this  ap- 
paratus the  joints  between  the  metal  shell  surrounding  the 
cell  are  made  by  stiff  rubber  packing,  hi  such  a  way  that  the 
osmotic  pressure  tends  to  tighten  the  connections.  An  ac- 
count of  the  construction  of  this  cell  will  soon  be  published. 

De  Vries  Method  of  Measuring  the  Relative  Osmotic 
Pressures  of  Solutions.  —  The  methods  thus  far  described 
measure  the  absolute  osmotic  pressures  of  solutions.  Be- 
fore leaving  the  subject  of  osmotic  pressure,  it  seems 
desirable  to  discuss  at  least  one  method  which  has  been 
described  for  measuring  the  relative  osmotic  pressures  ex- 
erted by  solutions  of  a  large  number  of  substances.  Hav- 
ing worked  out  the  relative  osmotic  pressures  of  these 
solutions,  and  the  absolute  osmotic  pressure  of  say  one  of 
them,  we  should  of  course  know  the  absolute  osmotic  pres- 
sures of  them  all. 

The  Dutch  botanist,  De  Vries,  was  the  first  to  develop 
and  apply  a  method  for  measuring  the  relative  osmotic 
pressures  of  solutions  of  a  number  of  electrolytes  and  non- 
electrolytes.  His  paper1  bore  the  title,  "Osmotic  Experi- 
ments with  Living  Membranes." 

The  method  of  De  Vries  consists  hi  preparing  solutions 
of  different  substances  of  such  concentrations,  that  they 
would  each  have  an  osmotic  pressure  just  equal  to  that 
of  a  given  solution  and,  therefore,  just  equal  to  one 
another.  The  solution  with  which  he  compared  the  solutions 
whose  relative  osmotic  pressures  he  was  studying,  was  the 
contents  of  the  cells  of  certain  plants;  and  here  the  difficulty 
arose  hi  finding  the  proper  plants.  We  can  see  what  this 
means  after  examining  the  method. 

i  Zett.  phys.  Chem.,  2,  415  (1888);  3, 103  (1889). 


68  THE  NATURE  OF  SOLUTION 

A  plant  cell  to  be  useful  for  this  work  must  consist  of  a 
solution  surrounded  by  a  semipermeable  membrane,  and 
this,  hi  turn,  surrounded  by  a  more  resistant  wall  of  cel- 
lulose to  prevent  the  membrane  from  being  easily  ruptured. 
When  such  a  cell  is  placed  hi  a  solution,  three  things 
may  happen,  depending  on  whether  the  osmotic  pressure 
of  the  solution  is  less  than,  greater  than,  or  just  equal  to 
that  of  the  cell  content.  If  the  osmotic  pressure  of  the 
solution  is  less  than  that  of  the  contents  of  the  cell,  water 
will  flow  from  the  solution  through  the  semipermeable  wall 
into  the  cell;  this  is  hi  accord  with  the  general  osmotic 
principle,  that  water  always  flows  from  the  region  of  lesser 
to  that  of  greater  osmotic  pressure.  The  cell  being  sur- 
rounded by  a  semipermeable  membrane  cannot  lose  its 
contents,  and  consequently  tends  to  become  distended. 

If  the  osmotic  pressure  of  the  solution  into  which  the 
cell  is  plunged  has  an  osmotic  pressure  which  is  greater 
than  that  of  the  cell  contents,  water  will  flow  from  the  cell 
outward  into  the  solution.  The  cell  having  lost  water,  its 
contents  will  shrink,  and  this  can  readily  be  observed  under 
the  microscope,  if  the  cell  fulfills  certain  conditions.  The 
third  and  last  possibility  is  where  the  osmotic  pressure  of 
the  solution  around  the  cell  is  the  same  as  that  of  the  cell 
content.  In  this  case  as  much  water  will  pass  into  the  cell 
hi  a  given  tune  as  will  pass  out,  and  the  cell  will  preserve 
its  normal  appearance. 

In  order  to  observe  this  change  hi  the  size  of  the  cell  its 
contents  must  be  colored,  and  here  a  difficulty  arose  in  finding 
the  proper  cells.  To  quote  the  words  of  De  Vries,  bearing 
on  this  point,  "Not1  every  plant  and  not  every  section  can 
be  employed.  We  need  cells  in  which  the  plasmolysis  can 
be  conveniently  observed,  and  a  section  hi  which  all  of  the 
cells  begin  to  show  this  phenomenon  at  exactly  the  same 
concentration  of  the  external  liquid."  The  only  plants 
which  at  that  tune  were  known  to  meet  these  requirements 
were:  Tradescantia  discolor,  Curcuma  rubricaulis  and  Begonia 

1  Zeit.  phys.  Chem.,  2,  417  (1888). 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  69 

manicata,  and  the  section  to  be  chosen  was  the  violet  to 
red  upper  layer  of  certain  parts  of  certain  leaves. 

Method  of  Procedure.  —  The  cells  were  placed  in  the 
solution  whose  osmotic  pressure  it  was  desired  to  compare 
with  that  of  the  cells.  They  were  observed  through  the 
microscope  to  see  whether  water  passed  into  or  out  of  the 
cell.  If  water  passed  into  the  cell,  shown  by  the  cell  becom- 
ing more  distended,  it  meant  that  the  osmotic  pressure  of 
the  solution  was  less  than  that  of  the  cell  contents.  In  this 
case  the  solution  was  made  more  and  more  concentrated, 
until  the  cells  when  plunged  into  it  remained  hi  a  perfectly 
normal  condition.  When  this  condition  was  reached,  it 
meant  that  the  osmotic  pressure  of  the  solution  was  just 
equal  to  that  of  the  contents  of  the  cell  —  the  two  were,  as 
we  say,  isosmotic. 

If,  on  the  other  hand,  the  cell  contents  shrank  when  the 
cell  was  placed  in  the  solution,  it  meant  that  the  osmotic 
pressure  of  the  solution  was  greater  than  that  of  the  cell 
content.  In  this  case  the  solution  was  diluted  step  by  step 
until  the  cell,  in  its  presence,  remained  hi  the  normal  state. 
When  this  was  realized,  it  meant  that  the  osmotic  pressures 
of  the  solution  and  of  the  cell  content  were  equal. 

These  solutions  which,  individually,  were  isosmotic  with 
the  contents  of  a  given  cell  or  kind  of  cells,  were  isosmotic 
with  one  another,  i.e.,  they  had  the  same  osmotic  pressures. 

Isotonic  Coefficients.  —  The  isosmotic  solutions  of  a 
large  number  of  different  substances  were  then  analyzed, 
and  their  concentrations  determined.  In  this  way  it  was 
learned  what  concentrations  of  different  substances  have  the 
same  osmotic  pressure  as  the  contents  of  a  given  cell,  and, 
consequently,  the  same  osmotic  pressure  as  one  another. 

De  Vries  expressed  the  concentrations  of  these  isosmotic 
solutions  not  hi  terms  of  gram-molecules  per  liter  of  solution, 
but  he  took  the  reciprocals  of  the  concentrations  and  called 
them  isotonic  coefficients.  De  Vries  worked  out  a  large  num- 
ber of  such  coefficients.  A  few  of  his  results  are  given  below.1 

1  Ibid.,  2,  425  (1888). 


70  THE  NATURE  OF  SOLUTION 


~  ,  Empirical  Isotonic 

Compound  tSSuto  Coefficient 


Glycerol      \  CsHaOa  1.78 

Cane  sugar  /  CuH.&On  1.81 

Malic  acid      \  C4H6O6  1.98 

Tartaric  acid  /  C4H6O«  2.02 

Potassium  nitrate    1  KNO,  3.0 

Sodium-  nitrate  NaNO3  3.0 

Potassium  chloride  f  KC1  3.0 

Sodium  chloride      \  NaCl  3.0 

Calcium  chloride       I  CaCl2  4.33 

Magnesium  chloride  /  MgCl2  4.33 

The  substances  that  are  bracketed  are  closely  related 
chemically,  and  it  will  be  observed  that  they  have  isotonic 
coefficients  of  the  same  order  of  magnitude. 

Limitations  of  the  Method.  —  It  is  evident  that  this 
method  is  limited  to  those  substances  and  classes  of  sub- 
stances which  do  not  act  chemically  upon  the  very  delicate 
semipermeable  membranes  which  surround  the  contents 
of  plant  cells.  This  means  that  the  method  is  limited 
practically  to  solutions  of  some  salts  and  neutral  organic 
compounds,  acids  and  bases  being  excluded.  However,  even 
with  this  limitation  the  method  has  given  results  of  value. 

Other  Methods  of  Determining  the  Relative  Osmotic 
Pressures  of  Solutions.  —  Hamburger1  found  that  he  could 
use  certain  cells  of  certain  animals  for  measuring  relative 
osmotic  pressures,  almost  as  well  as  De  Vries  had  used 
cells  of  plants.  The  red  blood  corpuscles  of  the  frog  and 
deer  were  adapted  to  this  purpose.  The  structure  of  these 
corpuscles  is,  from  our  present  standpoint,  a  solution  of  a 
large  number  of  things,  surrounded  by  a  semipermeable 
membrane  or  sac. 

The  principle  of  the  method  is  the  same  as  that  already 
described  with  the  vegetable  cells. 

Wladimiroff2  used  certain  bacteria  in  this  same  con- 
nection, determinhig  by  their  motions  whether  the  solution 
into  which  they  were  plunged  was  isosmotic  with  the  con- 
tents of  the  bacteria. 

1  Zeit.  phys.  Chem.,  6,  319  (1890). 

2  Ibid.,  7,  529  (1891). 


THE  OSMOTIC  PRESSURE  OF  SOLUTIONS  71 

These  methods  gave  results  comparable  with  those 
obtained  by  the  method  of  De  Vries. 

We  have  now  studied  the  methods  of  measuring  the  abso- 
lute osmotic  pressures  of  solutions,  and  the  methods  of  meas- 
uring their  relative  osmotic  pressures. 

We  must  next  ask,  why  was  osmotic  pressure  taken  up  hi 
the  present  connection?  What  bearing  have  the  measure- 
ments of  osmotic  pressure  on  the  nature  of  solution? 

We  shall  see  that  the  work  of  Pfeffer  was  fundamental 
to  the  recent  advances  which  have  been  made  in  our  knowl- 
edge of  matter  hi  the  dissolved  state.  This  will  be  devel- 
oped in  the  chapter  which  follows. 


CHAPTER  IV 

RELATIONS   BETWEEN    SOLUTIONS    AND    GASES 
DEMONSTRATED    BY    VAN'T   HOFF 

IN  the  historical  development  of  the  present  views  as 
to  the  nature  of  solution,  we  saw  that  Gay-Lussac,  as  early 
as  1839,  pointed  out  certain  relations  between  matter  hi  the 
dissolved  and  hi  the  gaseous  state  (p.  30).  A  similar  con- 
ception seems  to  have  been  held  a  little  later  (1845)  by  an 
Italian,  Bizio;  and  Rosenstiehl,  in  1870,  expressed  the 
opinion  that  there  is  an  analogy  between  the  osmotic  force 
which  exists  hi  solutions  and  the  elastic  force  hi  gases. 
Horstmann,  hi  1869,  found  that  the  equations  for  equi- 
librium in  gases  hold  for  solutions. 

We  have  already  become  familiar  with  the  views  of 
Guldberg  and  Waage  (p.  36),  Thomsen  (p.  42)  and  Men- 
del£eff  (p.  44),  on  the  analogy  between  solutions  and  gases. 

The  great  step,  however,  still  remained  to  be  taken, — 
the  mathematical  deduction  of  the  relation  between  matter 
hi  these  two  apparently  widely  different  states  of  aggrega- 
tion. This  was  done  by  the  Dutch  chemist,  Van't  Hoff. 

Van't  Hoff  published,  hi  1887,  an  epoch-making  paper1 
in  the  first  volume  of  the  "Zeitschrift  fur  physikalische 
Chemie"  on  "The  Role  of  Osmotic  Pressure  in  the  Analogy 
Between  Solutions  and  Gases."  This  title  will  indicate 
why  so  much  stress  was  laid  hi  the  last  chapter  on  the 
methods  employed  for  measuring  osmotic  pressure,  and 
especially  upon  the  results  that  have  been  obtained. 

Starting  with  Pfeffer's  results,  Van't  Hoff  pointed  out 
the  following  analogies,  and  deduced  the  following  relations. 

Said  Van't  Hoff,  if  there  is  any  fundamental  relation 
between  matter  in  the  dissolved  and  hi  the  gaseous  states, 

i  Zeit.  phys.  Chem.,  1,  481  (1887). 


RELATIONS  BETWEEN  SOLUTIONS  73 

then  the  generalizations  or  laws  which  hold  for  the  one 
condition  ought  to  apply  to  the  other.  The  best  known  and 
most  important  laws  of  gas  pressure  are  those  of  Boyle, 
Gay-Lussac,  and  Avogadro. 

The  law  of  Boyle  says  that  the  pressure  of  a  gas  is  pro- 
portional to  its  concentration. 

The  law  of  Gay-Lussac  states  that,  volume  remaining 
constant,  the  pressure  of  a  gas  increases  ^  for  every  rise  hi 
temperature  of  one  degree  centigrade;  or,  pressure  remaining 
constant,  the  volume  of  a  gas  increases  ^  for  every  rise 
in  temperature  of  one  degree. 

The  law  of  Avogadro,  as  usually  stated,  is  that  tem- 
perature and  pressure  being  constant,  equal  volumes  of 
all  gases  contain  the  same  number  of  ultimate  parts  or 
molecules.  For  our  present  purpose  it  may  best  be  stated 
as  follows:  Temperature  being  constant,  equal  volumes  of 
all  gases  which  contain  the  same  numbers  of  independent 
particles  exert  the  same  pressures. 

Law  of  Boyle  Applies  to  the  Osmotic  Pressures  of 
Dilute  Solutions.  —  From  the  osmotic  pressure  results, 
especially  of  Pfeffer,  Van't  Hoff  showed  that  the  law  of 
Boyle  for  gases  applies  to  the  osmotic  pressures  of  solutions. 
He  quoted  Pfeffer's  results,1  and  in  the  following  manner. 
If  we  represent  the  concentrations  of  the  solutions  by  C, 
and  the  osmotic  pressures  by  P,  we  have  the  following: 

r  P  * 

c 

1  %  535  mm.  535 

2%  1016  "  508 

2.74%  1518  "  554 

4%  2082  "  521 

6%  3075  "  513 

The  osmotic  pressure  divided  by  the  concentration  is  a 
constant  to  within  the  error  of  Pfeffer's  measurements. 
This  means,  of  course,  that  the  osmotic  pressure  is  propor- 
tional to  the  concentration,  which  is  simply  Boyle's  law  for 
gases. 

1  Ibid.,  1,  484  (1887). 


74  THE  NATURE  OF  SOLUTION 

Law  of  Gay-Lussac  Applies  to  the  Osmotic  Pressure  of 
Dilute  Solutions.  —  Van't  Hoff,  having  found  that  one  of 
the  laws  of  gas  pressure  applies  to  the  osmotic  pressures 
of  solutions,  went  farther  to  see  whether  other  laws  of 
gas  pressure  apply  to  solutions.  He  next  tested  the  applica- 
bility of  the  law  of  Gay-Lussac.1  He  did  this  both  theo- 
retically and  experimentally. 

It  will  be  recalled  that  Pfeffer  measured  not  only  the 
osmotic  pressures  of  a  number  of  solutions  of  several  sub- 
stances at  the  same  temperature,  but  also  the  osmotic  pres- 
sures of  a  given  solution  at  more  than  one  temperature.  He 
thus  furnished  data  for  calculating  the  temperature  coef- 
ficients of  osmotic  pressure.  A  solution  of  cane  sugar  which 
at  14.15°  exerted  an  osmotic  pressure  of  512  mm.,  at  32° 
gave  a  pressure  of  544  mm.;  while  a  solution  of  sodium 
tartrate,  which,  at  13.3°,  gave  a  pressure  of  1431.6  mm., 
at  36.6°  'showed  a  pressure  of  1564  mm. 

The  temperature  coefficients  of  osmotic  pressure,  per 
degree  rise  in  temperature,  calculated  from  the  data  fur- 
nished by  Pfeffer,  were  not  Yy^  as  would  be  expected  from 
the  law  of  Gay-Lussac,  but  were  of  this  order  of  magnitude. 
Were  we  dependent  upon  the  results  of  Pfeffer  alone  to 
determine  the  applicability  of  the  law  of  Gay-Lussac  to  the 
osmotic  pressures  of  solutions,  we  would  not  be  able  to  say 
whether  the  law  holds  rigidly  or  not;  but  fortunately  such 
is  not  the  case. 

Here  again  the  work  of  De  Vries  and  Hamburger  comes 
to  our  aid.  They  showed  that  solutions  which  were  pre- 
pared isosmotic  at  any  given  temperature  remained  isos- 
motic  when  the  temperature  was  varied.  This  shows  simply 
that  the  effect  of  temperature  on  the  osmotic  pressures  of 
the  different  solutions  is  the  same  —  the  different  solutions 
have  the  same  temperature  coefficients  of  osmotic  pressure. 
This  does  not  show  the  order  of  magnitude  of  the  tempera- 
ture coefficient,  and  therefore  does  not  show  whether  the  law 
of  Gay-Lussac  holds  for  the  osmotic  pressures  of  solutions. 

1  Zeit.  phys.  Chem.,  1,  486  (1887). 


RELATIONS  BETWEEN  SOLUTIONS  75 

While  we  are  not  able  to  decide  directly  the  question 
of  the  rigid  applicability  of  Gay-Lussac's  law  to  the  osmotic 
pressures  of  solutions,  we  can  decide  it  indirectly. 

Work  of  Soret.  —  In  1881  an  interesting  and  important 
paper  was  published  by  Soret,1  entitled,  "On  the  State  of 
Equilibrium  with  Reference  to  Concentration  of  a  Solid 
Solution  Originally  Homogeneous:  the  Two  Parts  having 
been  kept  at  Different  Temperatures."  This  was  a  continua- 
tion of  work  published  by  Soret  in  1879.  The  point  brought 
out  in  this  paper  is  that  a  homogeneous  solution  remains 
homogeneous  only  as  long  as  the  different  parts  are  kept 
at  the  same  temperature.  When  one  part  of  an  initially 
homogeneous  solution  is  kept  at  a  different  temperature 
from  the  other,  the  solution  becomes  heterogeneous: 
the  colder  part  becoming  the  more  concentrated,  and  the 
warmer  more  dilute.  This  is  the  "  Principle  of  Soret." 
The  question  arises,  what  causes  this  difference  in  concen- 
tration? and  the  answer  is  obvious,  diffusion.  The  dis- 
solved substance  diffuses  from  the  warmer  into  the  colder 
part  of  the  solution.  We  shall  see  later  that  all  diffusion 
in  solution  is  caused  by  osmotic  pressure,  and  therefore 
the  principle  of  Soret  enables  us  to  test  the  applicability 
of  the  law  of  Gay-Lussac  to  the  osmotic  pressures  of 
solutions. 

If  this  law  holds,  then  the  difference  in  concentration 
between  the  colder  and  warmer  parts  of  the  solution  must 
be  STF  of  the  original  concentration  for  each  degree  dif- 
ference in  temperature.  What  was  the  difference  found 
experimentally? 

In  his  earlier  experiments,  where  the  different  parts  of 
the  solution  were  allowed  to  stand  at  the  two  temperatures 
only  twenty-three  days,  Soret  found  differences  in  con- 
centration which  were  always  less  than  would  be  calculated 
from  the  law  of  Gay-Lussac.  In  his  later  work,2  where  the 
tubes  containing  the  solution  were  allowed  to  stand,  the  two 
parts  at  the  different  temperatures,  for  from  fifty  to  fifty- 

1  Ann.  Chim.  Phys.  [5],  22,  293  (1881).  *  Loc.  tit. 


76  THE  NATURE  OF  SOLUTION 

six  days,  the  differences  in  the  concentrations  of  the  solution 
in  the  warmer  and  in  the  colder  portions  were  greater  than 
in  his  earlier  experiments,  showing  that  in  the  earlier  work 
the  solution  had  not  stood  long  enough  at  the  two  tem- 
peratures to  establish  equilibrium. 

From  Soret's  later  results  Van't  Hoff  made  the  follow- 
ing calculations.  One  part  of  a  solution  was  cooled  to  20° 
and  the  other  warmed  to  80°.  The  concentration  of  the 
cooled  portion  was  17.33  percent.  Assuming  Gay-Lus- 
sac's  law  to  hold,  it  was  calculated  that  the  concentration 
of  the  warmer  part  of  the  solution  should  be  14.3  percent; 
while  the  concentration  found  was  14.03  percent. 

In  another  experiment  the  concentration  in  the  part 
of  the  solution  which  was  at  20°  was  29.87  percent.  From 
this  it  was  calculated  that  the  concentration  of  the  part 
warmed  to  80°  should  be  24.8  percent;  while  the  value 
found  was  23.87  percent. 

It  will  be  observed  that  the  concentration  calculated 
for  the  warmer  solution  was  always  slightly  greater  than 
that  found  experimentally,  even  when  the  solution  had  stood, 
the  two  parts  at  the  different  temperatures,  for  fifty  days. 

It  has  since  been  calculated  that  even  fifty  days  is  not 
long  enough,  under  the  conditions  of  Soret's  experiments, 
for  the  tubes  containing  the  solution  to  stand  hi  order 
that  equilibrium  might  be  reached.  Taking  this  into  ac- 
count, and  the  comparatively  small  difference  noted  above 
between  the  values  found  and  those  calculated,  we  seem 
reasonably  safe  in  concluding  that  the  law  of  Gay-Lussac 
for  gases,  like  the  law  of  Boyle,  applies  to  the  osmotic  pres- 
sure of  solutions. 

Equality  of  Gas-Pressure  and  of  Osmotic  Pressure. — 
We  have  shown  that  gas-pressure  is  proportional  to  osmotic 
pressure.  The  applicability  of  the  two  laws  of  gases  just 
considered  establishes  this  relation.  What  we  want  to 
know,  however,  is  the  relation  between  the  actual  osmotic 
pressure  and  the  actual  gas  pressure  when  the  gas  and  the 
solution  are  of  the  same  concentration,  i.e.,  contain  the  same 


RELATIONS  BETWEEN  SOLUTIONS  77 

number  of  parts  in  a  given  volume  —  does  Avogadro's  law 
for  gases  apply  to  solution? 

Van't  Hoff  has  furnished  this  information. 

The  way  to  test  this  point  is  to  prepare  a  solution  of 
known  concentration,  i.e.,  containing  a  certain  number  of 
dissolved  parts  in  a  certain  volume  of  the  solution;  and 
determine  the  osmotic  pressure  of  this  solution.  Then 
take  a  gas  of  such  concentration  that  it  contains  the  same 
number  of  gas-particles  hi  the  same  volume.  Measure  the 
pressure  of  this  gas.  Then  compare  directly  the  osmotic 
pressure  of  the  solution  with  the  gas-pressure  of  the  gas. 
This  comparison  was  made  by  Van't  Hoff.1 

He  took  a  one  percent  solution  of  cane  sugar,  i.e.,  a 
solution  containing  one  gram  of  sugar  hi  one  hundred  grams 
of  water.  This  solution,  considered  from  the  standpoint 
of  volume,  contains  a  gram  of  sugar  hi  100.6  c.c.  of  solu- 
tion. From  Pfeffer's  data  Van't  Hoff  calculated  the  osmotic 
pressure  exerted  by  such  a  solution. 

If  we  compare  this  with  the  gas-pressure  exerted  by  a 
quantity  of  gas,  say  hydrogen,  containing  the  same  number 
of  gas  particles  as  there  were  sugar  particles  hi  the  same 
volume,  we  would  have  the  following: 

Temperature  (t)  Osmotic  pressure  Gas  pressure 

6.8°  0.664  mm.  0.665  mm. 

13.7  0.691  0.681 

14.2  0.671  0.682 

15.5  0.684  0.686 

22.0  0.721  0.701 

32.0  0.716  0.725 

36.0  0.746  0.735 

The  above  data  show  that  the  two  pressures  are  approxi- 
mately equal  to  one  another,  under  the  same  conditions 
of  concentration  and  temperature. 

A  moment's  thought  will  show  that  this  equality  is  very 
surprising.  Think,  on  the  one  hand,  of  a  gas.  The  parti- 
cles are  free  to  move,  and  do  move  with  very  high  velocities, 
exerting  a  pressure  due  to  their  impacts  against  the  walls  of 
the  containing  vessel. 

Think,  on  the  other  hand,  of  a  solution.    The  dissolved 

1  Zeit  phys.  Chem.,  1,  493  (1887). 


78  THE  NATURE  OF  SOLUTION 

particles  move  very  slowly  through  the  viscous  liquid. 
The  osmotic  pressure  which  they  exert,  it  seems,  can  never 
be  accounted  for  on  a  purely  kinetic  basis.  Indeed,  we  do 
not  know  today  what  is  the  cause  of  osmotic  pressure. 
Yet,  osmotic  pressure  hi  result,  i.e.,  hi  order  of  magnitude, 
is  equal  to  gas-pressure.  This  shows,  of  course,  since 
Avogadro's  law  holds  for  the  gas-pressures  of  gases,  that  it 
must  hold  for  the  osmotic  pressures  of  solutions. 

Van't  Hoff1  also  deduced  this  same  relation  on  purely 
theoretical  grounds. 

Importance  of  the  Applicability  of  the  Gas-Laws  to 
Solutions.  —  We  must  now  ask  what  has  been  gained  by 
showing  that  the  laws  of  gas-pressure  apply  to  the  osmotic 
pressures  of  solutions?  How  does  this  aid  us  hi  dealing 
with  the  problem  of  solutions? 

About  matter  in  the  solid  state  we  really  know  very 
little.  We  know  the  shapes  of  the  crystals  which  it  forms. 
We  know  a  little  about  the  passage  of  the  various  forms  of 
energy  —  light,  heat,  electricity  —  through  solids;  but  this 
knowledge  is  after  all  superficial.  Our  ignorance  of  the 
real  nature  of  solids  is  very  nearly  perfect. 

When  we  turn  to  liquids,  we  know  very  much  more 
about  them  than  we  do  about  solids.  We  know  the  molecular 
weights  of  substances  in  the  liquid  state;  we  have  studied 
pretty  thoroughly  the  physical  properties  of  liquids.  In  a 
word,  we  know  a  good  deal  about  liquids;  but  even  here 
what  we  know  is  comparatively  little  hi  comparison  with 
what  we  do  not  know. 

When  we  turn  to  gases,  we  really  know  much  more. 
Here  is  matter  in  the  most  dilute  and  in  the  simplest  state. 
The  laws  of  gases  are  comparatively  simple,  and  are  pretty 
well  understood.  Matter  in  the  gaseous  state  lends  itself 
to  mathematical  treatment  as  matter  in  no  other  state  of 
aggregation  does,  and  this  is  the  most  important  point.  We 
can  apply  the  only  rigid  and  exact  scientific  method  — 
the  mathematical  —  to  gases. 

1  Zeit.  phys.  Chem.,  1,  489  (1887). 


RELATIONS  BETWEEN  SOLUTIONS  79 

Van't  Hoff  in  this  work,  then,  has  shown  that  we  can  deal 
with  solutions  largely  as  we  deal  with  gases  —  since  they 
both  obey  the  same  fundamental  laws;  and  this  is  a  great 
step  forward  hi  our  dealing  with  solutions.  We  can  now 
deal  with  them,  to  a  very  considerable  extent,  by  the  exact 
mathematical  method,  and  thereby  place  our  knowledge 
of  dissolved  substances  upon  a  more  scientific  basis. 

In  the  first  chapter  of  this  work,  attention  was  called  to 
the  importance  of  solution  for  science  La  general,  and  for 
chemistry  in  particular.  We  saw  that  it  is  absolutely 
essential  for  chemistry  —  chemistry  is  a  branch  of  the 
science  of  solutions.  Now  that  we  can  deal  with  solution 
by  the  truly  scientific  method  —  the  mathematical  —  and 
since  solution  lies  right  at  the  foundation  of  all  chemistry, 
it  is  obvious  that  we  have  made  good  progress  hi  the 
direction  of  converting  chemistry  into  a  branch  of  exact 
science,  like  the  sister  science,  physics.  Much  has  already 
been  accomplished  hi  this  direction.  Chemistry  is  becom- 
ing more  and  more  mathematical,  and  will  become  still 
more  mathematical  hi  the  near  future.  This  is  due  hi  no 
small  degree  to  the  relations  pointed  out  by  Van't  Hoff 
between  solutions  and  gases. 

The  importance  of  this  generalization,  mathematically 
deduced,  cannot  be  easily  overestimated.  If  the  author 
were  asked  to  select  the  most  important  generalization 
which  has  ever  been  reached  hi  chemistry,  tending  to  con- 
vert chemistry  from  empiricism  into  science,  he  would 
unhesitatingly  name  the  application  of  the  gas  laws  to 
solution. 

Although  this  was  hinted  at  in  a  qualitative  way  by 
several  before  the  time  of  Van't  Hoff,  it  remained  for  this 
great  Dutch  chemist,  at  the  age  of  thirty-five,  to  deduce 
these  relations  mathematically,  and  thus  place  them  once 
for  all  upon  a  scientific  basis. 

Exceptions  to  the  Above  Relations. —  Van't  Hoff  has, 
then,  shown  that  the  three  fundamental  laws  of  gas-pressure 
—  the  laws  of  Boyle,  Gay-Lussac,  and  Avogadro  —  apply  to 


80  THE  NATURE  OF  SOLUTION 

the  osmotic  pressures  of  solutions.  While  this  is  true,  it  is 
not  the  whole  truth.  These  laws  apply  to  the  osmotic 
pressures  of  solutions  of  certain  kinds  of  substances,  but 
they  do  not  apply  at  all  to  the  osmotic  pressures  of  solutions 
of  other  types  of  compounds. 

These  laws  hold  for  solutions  of  nonelectrolytes,  i.e., 
solutions  of  substances  which  do  not  conduct  the  current. 
These  include  all  of  the  neutral  organic  compounds;  in  a 
word,  all  substances  except  acids,  bases,  and  salts.  The 
gas  laws  do  not  apply  to  the  osmotic  pressure  of  a  single 
electrolyte;  and  the  electrolytes,  acids,  bases,  and  salts, 
are  by  far  the  most  interesting  and  important  chemical 
compounds.  They  are  the  active  substances,  the  things 
that  primarily  give  us  chemistry.  The  nonelectrolytes  are 
comparatively  inactive. 

Van't  Hoff  called  attention  to  the  fact  that  the  electro- 
lytes exert  too  great  osmotic  pressure,  in  terms  of  that  exerted 
by  the  nonelectrolytes,  which,  we  have  seen,  was  exactly 
equal  to  that  exerted  by  gases  at  the  same  concentration  as 
the  solution  of  the  nonelectrolyte. 

Even  Van't  Hoff  could  not  explain  this  discrepancy 
presented  by  the  electrolytes.  This  remained  for  another  — 
Svante  Arrhenius. 


CHAPTER  V 

THE   THEORY    OF   ELECTROLYTIC    DISSOCIATION    AS 
ANNOUNCED    BY    ARRHENIUS 

IN  Chapter  II  of  this  work,  we  saw  that  a  number  of 
investigators  had  in  mind  the  conception  that  in  solution 
molecules  of  substances  are  in  some  way  broken  down  into 
parts. 

Williamson  (p.  31)  was  forced  to  this  conclusion  from 
his  study  of  the  action  of  sulphuric  acid  on  alcohol. 

Clausius  (p.  33),  from  the  study  of  the  electrical  be- 
havior of  solutions  of  electrolytes,  concluded  that  such 
solutions  must  contain  not  only  whole  molecules,  but 
also  part  molecules. 

From  the  additive  property  of  the  constituents  of  salt 
solutions  Valson  (p.  36)  thought  that  the  constituents 
of  the  salt  hi  solution  must  be  in  a  state  of  virtual  inde- 
pendence, and  Favre  and  Valson  (p.  38)  were  of  the  same 
opinion. 

Kohlrausch  (p.  40)  discovered  the  law  of  the  independent 
migration  velocities  of  the  ions  —  another  additive  property 
of  solutions,  and  Raoult  (p.  43)  added  other  new  proper- 
ties of  solutions  —  lowering  of  freezing-point  and  lowering 
of  vapor-tension  —  to  those  which  were  the  sum  of  the  cor- 
responding properties  of  the  constituents  of  the  dissolved 
substances. 

The  Italian,  Bartholi,  in  1882,  seemed  to  have  had  in 
mind  some  conception  of  dissociation  of  the  dissolved  parts 
hi  solution. 

It  remained,  however,  for  the  young  Swedish  physicist, 
Svante  Arrhenius,  to  give  us  the  theory  of  electrolytic 
dissociation  hi  the  form  in  which  we  now  have  it,  and 
which  has  proved  to  be  one  of  the  corner  stones  in  the 


82  THE  NATURE  OF  SOLUTION 

epoch-making  developments  in  chemistry  during  the  last 
thirty  years. 

Arrhenius  and  the  Dissociation  Theory.  —  Just  at  the 
tune  (1887)  that  Van't  Hoff  encountered  the  apparent  discrep- 
ancies presented  by  electrolytes,  hi  regard  to  the  relations 
between  solutions  and  gases,  this  young  Swede  came  to  his 
laboratory.  In  1883  Arrhenius  had  studied  the  electrical 
conductivity  of  solutions  of  electrolytes,  and  hi  1884  had 
published  the  view  that  such  solutions  contain  two  kinds  of 
molecules,  those  that  conduct  because  they  are  decom- 
posed into  their  part  molecules,  and  those  that  do  not  con- 
duct because  they  are  present  as  whole  molecules.  With 
this  conception  hi  mind,  Arrhenius  turned  to  the  problem 
of  the  meaning  of  the  abnormally  large  osmotic  pressures 
exerted  by  electrolytes,  which  do  not  obey  the  laws  of  gas- 
pressure,  and  which  compelled  Van't  Hoff  to  introduce  into 
the  simple  gas  equation  PV  =RT,  the  coefficient  i,  which 
was  always  greater  than  unity,  hi  order  that  the  equation 
might  be  applied  to  the  osmotic  pressures  of  such  solutions. 
What  is  the  significance  of  these  abnormally  great  osmotic 
pressures? 

The  problem  seems  to  have  presented  itself  to  Arrhenius' 
mind  about  as  follows.  The  osmotic  pressures  of  solu- 
tions of  nonelectrolytes  obey  the  law  of  Boyle.  This  sim- 
ply means  that  these  osmotic  pressures  are  a  function  of 
numbers  —  of  the  ratio  between  the  number  of  parts  of  the 
dissolved  substance  and  the  number  of  parts  of  the  solvent. 
This  is  what  might  be  called  an  arithmetical  property. 

Since  osmotic  pressure  is  a  function  of  numbers  and  num- 
bers only,  when  we  have  too  great  osmotic  pressure  we 
must  have  a  too  great  number  of  parts  present.  This  con- 
clusion is  unavoidable.  The  solution  of  this  problem  to 
a  mathematical-physical  mind  like  Arrhenius',  was  obvious. 
If  we  want  to  account  for  a  larger  number  of  parts  than 
correspond  to  the  whole  molecules,  we  must  consider  these 
whole  molecules  as  broken  down  into  such  parts,  and  this 
is  what  Arrhenius  did. 


THEORY  OF  ELECTROLYTIC  DISSOCIATION  83 

He  assumed  that  molecules  of  acids,  bases,  and  salts  in 
solution  are  broken  down,  or  dissociated,  into  parts;  these 
parts  being  shown  to  be  charged  because  they  conduct 
the  current.  Arrhenius,  however,  went  much  farther  than 
this.  He  did  not  simply  say,  as  did  Clausius,  that  same  of 
the  molecules  of  an  electrolyte  are  broken  down  into  part 
molecules,  but  hi  his  epoch-making  paper  published  hi  the 
first  volume  of  the  "Zeitschrift  fur  physikalische  Chemie,"1 
he  points  out  methods  for  determining  the  relative  numbers 
of  the  molecules  that  are  broken  down  or  dissociated 
into  part  molecules. 

Arrhenius  Points  out  Methods  of  Measuring  Disso- 
ciation.2 —  As  has  been  stated,  as  early  as  1883  Arrhenius 
was  working  on  the  property  of  aqueous  solutions  of  elec- 
trolytes to  conduct  the  current.  The  part  molecules  are 
those  which  carry  the  current,  and  these  he  termed  the 
"  active"  molecules.  Those  molecules  which  do  not  take 
part  in  conducting  the  current  he  termed  "inactive."  The 
inactive  whole  molecules  break  up  into  the  active  part 
molecules,  called  ions,  which  are  free  to  move  independent 
of  one  another.  In  the  inactive  molecules  the  constit- 
uents are  bound  to  one  another.  As  we  increase  the 
dilution  of  a  solution,  the  inactive  molecules  pass  over 
into  active,  until  at  great  dilution  practically  all  of  the 
molecules  have  become  active.  Arrhenius  calls  the  ratio 
between  the  number  of  active,  and  the  sum  of  the  active 
and  inactive  molecules  the  "  activity  coefficient."  When  all 
of  the  inactive  have  passed  over  into  active  molecules, 
this  coefficient  is,  of  course,  unity;  for  smaller  dilutions 
less  than  unity.  For  any  dilution  the  activity  coefficient, 
which  is  what  we  today  call  the  percentage  dissociation  a, 
is  the  ratio  between  the  molecular  conductivity  at  the 
dilution  in  question  (juc),  and  the  molecular  conductivity 
infinite  dilution  (MOO). 


Zeit.  phys.  Chem.,  1,  631  (1887).  •  Ibid.,  p.  633. 


84  THE  NATURE  OF  SOLUTION 

Knowing  the  activity  coefficient,  we  can  calculate  the 
value  of  i  in  the  Van't  Hoff  equation  for  osmotic  pressure. 
If  we  represent  the  number  of  ions  into  which  every  molecule 
breaks  down  by  kf 

i  =  1  +  (k  - 1)  a 

The  coefficient  i  can  also  be  calculated  from  the  lowering 
of  the  freezing-point,  i  is  the  ratio  between  the  molecular 
lowering  of  the  freezing-point  observed,  and  the  molecular 
lowering  if  there  were  no  dissociation.  Let  us  represent  the 
former  by  £,  the  latter  is  18.5  for  water;  the  coefficient  i 
is  calculated  thus: 

. t_ 

"18.5 

The  terms  molecular  conductivity  and  molecular  lower- 
ing of  the  freezing-point  have  recently  been  used.  Molec- 
ular conductivity  means  the  actual  conductivity  divided 
by  the  concentration  expressed  decimally,  and  molecular 
lowering  of  the  freezing-point  means  the  actual  lowering  of 
the  freezing-point  divided  by  the  concentration  expressed 
decimally. 

Dissociation  from  Freezing-Point  Lowering  Compared 
with  Dissociation  from  Conductivity.  —  Probably  the  most 
striking  feature  of  Arrhenius'  paper  is  his  comparison 
of  electrolytic  dissociation  as  calculated  from  freezing- 
point  measurements,  with  the  results  obtained  from  the 
study  of  the  conductivity  of  solutions.  A  few  of  his 
results  for  nonelectrolytes,  and  for  acids,  bases,  and  salts 
are  given  below.1 

NONELECTROLYTES 

Compound  Formula  a  i  =:— ~r  i  =  l  +  (fc-l)a 

18.5 

Methyl  alcohol  CH4O  0  0.94  .00 

Glycerol  C3H5(OH)3  0  0.92  .00 

Cane  sugar  C12H.>2Oii  0  1.00  .00 

Acetone  C3H6O  0  0.92  .00 

Ethyl  acetate  C4H8O2  0  0.96  .00 

Acetamide  C2H3ONH2  0  0.96  .00 

1  Zeit.  phys.  Chem.,  1,  634  (1887). 


THEORY  OF  ELECTROLYTIC  DISSOCIATION 


85 


Acros 


Compound 


Formula 


Hydrochloric  acid 
Nitric  acid 
Sulphuric  acid 
Hydrogen  sulphide 
Hydrocyanic  acid 
Acetic  acid 
Malic  acid 

HC1 
HNO3 

HCN 
CH3COOH 

0.90 
0.92 
0.60 
0.00 
0.00 
0.01 
0.04 


Compound 

Sodium  hydroxide 
Potassium  hydroxide 
Calcium  hydroxide 
Strontium  hydroxide 
Barium  hydroxide 
Ammonia 
Methvlamine 
Ethylamine 


Compound 

Potassium  chloride 
Sodium  chloride 
Potassium  bromide 
Sodium  nitrate 
Sodium  acetate 
Potassium  chlorate 
Sodium  carbonate 
Ammonium  sulphate 
Barium  nitrate 
Strontium  nitrate 
Lead  nitrate 
Cupric  acetate 
Mercuric  chloride 
Cadmium  nitrate 


BASES 

Formula 

a 

NaOH 

0.88 

KOH 

0.93 

Ca(OH)2 

0.80 

Sr(OH)2 

0.86 

Ba(OH)2 

0.84 

NH3 

0.01 

CH3NH2 

0.03 

CzHsNH, 

0.04 

SALTS 

)rmula 

a 

KCl 

0.86 

NaCl 

0.82 

KBr 

0.92 

NaNO« 

0.82 

CHsCOONa 

0.79 

KC1O3 

0.83 

NaoCO3 

0.61 

(NH4)2SO4 

0.59 

Ba(N03)2 

0.57 

Sr(N03)2 

0.62 

Pb(N03)2 

0.54 

(CiH,0»)iCu 

0.33 

HgCl2 

0.03 

Cd(N03)2 

0.73 

=    t 

18.5 
1.98 
1.94 
2.06 
1.04 
1.05 
1.01 
1.08 


18.5 
1.96 
1.91 
2.59 
2.61 
2.69 
1.03 
1.00 
1.00 


18.5 
1.82 
1.90 
1.90 
1.82 
1.73 
1.78 
2.18 
2.00 
2.19 
2.23 
2.02 
1.68 
1.11 
2.32 


t  -  1  +  (*  -  1)  a 

1.90 
1.92 
2.19 
1.00 
1.00 
1.01 
1.07 


1.88 
1.93 
2.59 
2.72 
2.67 
1.01 
1.03 
1.04 


1.86 
1.82 
1.92 
1.82 
1.79 
1.83 
2.22 
2.17 
2.13 
2.23 
2.08 
1.66 
1.05 
2.46 


The  importance  of  these  quantitive  results  is  very  great 
indeed.  The  agreements  are  far  too  close  and  exist  in  such 
a  large  number  of  cases,  that  it  is  impossible  to  account  for 
them  as  due  to  accident.  The  impression  that  these  rela- 
tions made  upon  the  scientific  world  was  very  deep.  There 
were  agreements  that  could  not  be  ignored,  and  although 
the  assumption  upon  which  they  were  calculated  was  quite 
revolutionary  hi  character,  it  seemed  from  these  agreements 
that  it  must  contain  a  large  element  of  truth.  These 


86  THE  NATURE  OF  SOLUTION 

results  undoubtedly  did  more  to  attract  men  of  science 
to  the  Arrhenius'  theory  than  any  other  one  argument 
bearing  upon  it. 

Arrhenius,  hi  discussing  these  relations,  calls  attention 
to  the  fact  that  in  making  these  calculations  he  has  made 
two  assumptions:  viz.,  that  the  law  of  Van't  Hoff,  or  the 
relations  between  the  osmotic  pressure  of  solutions  and  the 
gas-pressure  of  gases,  holds  not  simply  for  nonelectrolytes, 
but  also  for  the  electrolytes,  provided  we  take  into  account 
their  degree  of  dissociation;  and,  secondly,  that  an  aqueous 
solution  of  any  electrolyte  is  partly  dissociated,  i.e.,  consists 
partly  of  undissociated  molecules  and  partly  of  ions,  the 
relative  numbers  of  the  two  being  a  function  of  the  dilution 
of  the  solution.  As  the  dilution  of  the  solution  is  increased 
the  number  of  active  molecules  increases,  the  dissociation 
becoming  complete  at  infinite  dilution. 

Arrhenius  then  takes  up  a  discussion  of  certain  proper- 
ties of  solutions  which  are  additive,  i.e.,  they  are  the  sum 
of  the  corresponding  properties  of  the  ions  of  the  solution. 
With  certain  of  these  properties  we  are  not  yet  supposed  to 
be  familiar,  and  they  will  therefore  not  be  discussed  in  this 
connection. 

Bearing  of  the  Theory  of  Electrolytic  Dissociation  on 
Chemistry.  —  The  theory  of  electrolytic  dissociation,  as  it 
was  given  us  by  Arrhenius,  says  that  dilute  solutions  of 
acids,  bases,  and  salts  contain  practically  ions  only,  and 
very  few  molecules.  It  will  be  recognized  that  ions  are 
the  substances  which  are  the  most  active  chemically.  If 
dilute  solutions  of  these  substances  contain  only  ions  and 
practically  no  molecules,  then  the  chemical  activity  of  these 
solutions  cannot  be  due  to  molecules,  since  there  are  but 
few  molecules  present,  but  must  be  due  to  the  ions.  It 
was  further  shown  that  dilute  solutions  of  acids,  bases  and 
salts  are  more  active,  if  we  take  into  account  the  dilution, 
i.e.,  refer  everything  to  molecular  quantities,  than  more 
concentrated  solutions.  For  example,  a  thousandth  nor- 
mal solution  of  an  acid  is  more  than  one-thousandth  as 


THEORY  OF  ELECTROLYTIC    DISSOCIATION  87 

active  as  a  normal  solution.  This  would  indicate  not  only 
that  ions  are  chemically  active,  but  that  they  are  the  most 
active  chemical  agents. 

This  conclusion  at  the  tune  we  are  considering  was 
revolutionary.  Chemists  prior  to  this  time  had  always 
looked  upon  atoms  and  molecules  as  the  active  agents  of 
chemistry.  Indeed,  atomic  and  molecular  chemistry  was 
well-nigh  universally  accepted.  Chemists  had  become  ac- 
customed to  think  of  reactions  hi  terms  of  molecules  and 
atoms,  and  now  to  ask  them  to  lay  all  this  aside,  and  start 
over  again,  thinking  only  of  ions  or  charged  parts  as  ef- 
fecting chemical  reactions,  was  asking  a  good  deal.  Indeed, 
it  was  soon  seen  that  not  only  are  ions  the  most  active 
agents  chemically,  but  facts  were  brought  to  light  which 
made  it  highly  probable  that  ions  are  the  only  active  chemical 
agents.  Ions  were  apparently  necessary  in  order  that  we 
might  have  a  chemical  reaction.  Ions  could  react  with 
molecules  and  with  other  ions,  but  molecules  could  not 
react  with  other  molecules.  Some  facts  bearing  on  this 
point  have  been  discussed  hi  the  first  chapter  of  this  book. 

This  was  going  very  far  indeed.  To  ask  chemists  to 
add  to  atomic  and  molecular  chemistry,  ionic  chemistry, 
would  be  asking  much;  but  to  ask  them  to  abandon  atomic 
and  molecular  chemistry  and  to  substitute  for  it  ionic 
chemistry  was  sure  to  arouse  opposition,  especially  on  the 
part  of  those  who  had  been  accustomed  for  a  long  tune  to 
think  and  work  hi  terms  of  atoms  and  molecules. 

Opposition  to  the  Theory  of  Electrolytic  Dissociation. — 
Almost  as  soon  as  Arrhenius  proposed  the  dissociation 
theory  a  storm,  indeed  in  certain  quarters  a  veritable 
cyclone,  of  opposition  to  the  theory  broke  loose.  In  some 
cases  this  opposition  was  based  on  a  careful  study  of  the 
facts,  and,  in  other  cases,  on  an  inability  or  an  indisposition 
to  adapt  one's  self  to  the  new  ideas. 

A  careful  examination  will  show  that  hi  some  cases  the 
opposition  to  the  theory  of  electrolytic  dissociation  was 
based  in  no  small  measure  upon  a  certain  lack  of  familiarity 


88  THE  NATURE  OF  SOLUTION 

with  the  subject  that  was  being  discussed.  For  example, 
it  was  said  that  the  theory  of  electrolytic  dissociation  could 
not  be  true,  because  metallic  sodium  and  metallic  potassium 
could  not  remain  in  the  presence  of  water  without  acting 
chemically  upon  it.  This,  of  course,  entirely  failed  to  distin- 
guish between  an  atom  and  an  ion.  An  atom  is  uncharged, 
or  electrically  neutral;  an  ion  is  charged.  There  is  no 
reason  for  judging  of  the  properties  of  a  charged  atom  or 
group  of  atoms,  from  the  corresponding  properties  of  the 
same  atom  or  group  when  uncharged;  and  this  kind  of 
criticism,  therefore,  had  no  scientific  significance. 

This  example  is  cited  to  illustrate  a  certain  type  of 
criticism  to  which  the  theory  of  electrolytic  dissociation  was 
subjected  in  its  early  days.  It  is  obvious  that  this  type 
was  based  largely  upon  a  misconception  of  the  meaning  and 
of  the  significance  of  the  theory  of  electrolytic  dissociation. 

There  was,  on  the  other  hand,  a  more  conservative  criti- 
cism of  the  theory,  on  the  part  of  many  who  had  no  case  to 
advocate,  no  preconception  as  to  the  truth  in  the  matter,  no 
strong  prejudice  to  overcome.  They  were  prompted  only 
by  the  desire  to  get  at  the  truth;  and  the  new  theory  was 
called  upon  to  prove  its  case.  This  kind  of  criticism  is 
always  wholesome.  It  is  not  obstructively  conservative, 
but  the  new  theory  or  generalization  which  would  supplant 
existing  views,  is  called  upon  to  prove  itself.  This  is  exactly 
right,  and  it  is  this  kind  of  criticism  that  gets  at  the  truth. 
Those  critics  of  the  dissociation  theory  who  belonged  in 
this  class  studied  the  new  theory  in  as  many  of  its  bearings 
as  possible.  They  looked  into  the  facts  underlying  the 
theory  to  see  on  what  kind  of  a  foundation  it  rested,  and 
then  proceeded  to  ask  questions. 

They  made  one  point  which  was  at  that  tune  absolutely 
unanswerable. 

Relations  Between  Solutions  and  Gases  and  the  Theory 
of  Electrolytic  Dissociation  Hold  Only  for  Dilute  Solutions. 
—  Attention  was  called  to  the  fact  that  the  relations  pointed 
out  by  Van't  Hoff  between  solutions  and  gases,  as  well  as  the 


THEORY  OF  ELECTROLYTIC  DISSOCIATION  89 

theory  of  electrolytic  dissociation  itself,  hold  only  for  dilute, 
and  indeed,  only  for  very  dilute  solutions.  These  relations 
hold  rigidly  only  for  what  was  termed  "ideal  solutions." 
Said  the  critics  of  the  dissociation  theory,  this  may  all  be 
very  well  as  far  as  it  goes,  but  it  does  not  go  far  enough. 
We  want  a  theory  of  solution  which  does  not  apply  simply 
to  very  dilute  or  "  ideal  solutions/'  but  which  holds  for  all 
solutions.  It  was  pointed  out  that  these  relations  hold  only 
for  solutions  so  dilute  that  they  have  little  or  no  significance 
hi  chemistry.  They  do  not  apply  to  the  ordinary  solutions 
that  we  use  hi  the  chemical,  physical,  biological,  or  any 
other  scientific  laboratory. 

It  was  insisted,  and  rightly  so,  that  this  is  a  serious  defect 
in  the  dissociation  theory  and  hi  the  relations  between  the 
osmotic  pressures  of  solutions  and  the  gas  pressures  of  gases. 
These  relations  and  this  theory  of  solution  which  was  the 
outcome  of  these  relations  bear  somewhat  the  same  relation 
to  a  comprehensive  theory  of  solution  as  the  laws  of  Boyle 
and  Gay-Lussac  do  to  a  comprehensive  law  of  gas-pressure 
which  would  apply  under  a  great  variety  of  conditions,  such 
as  the  equation  of  Van  der  Waals. 

This  objection  was  not  answered,  and  hi  the  light  of  the 
facts  then  known  could  not  be.  It  was  a  weak  point  hi  the 
relations  between  solutions  and  gases  demonstrated  by 
Van't  Hoff,  and  in  the  theory  of  electrolytic  dissociation  as 
proposed  by  Arrhenius. 

It  seems  that  we  now  have  the  partial  explanation  of  why 
it  is  that  these  relations  and  this  theory  did  not  appear  to 
hold  for  the  more  concentrated  solutions.  This  will  be  dis- 
cussed at  some  length  in  the  last  two  chapters  of  this  work. 


CHAPTER  VI 

DIFFUSION   IN    SOLUTION 

The  Phenomenon  of  Diffusion.  —  When  a  solution  of  a 
given  concentration  is  brought  in  contact  with  the  pure 
solvent,  or  with  a  solution  of  the  same  substance  at  a  dif- 
ferent concentration,  the  dissolved  substance  passes  over 
from  the  more  concentrated  to  the  more  dilute  solution, 
and  continues  to  do  so  until  equality  of  concentration  is 
established. 

Place  a  solid  in  the  bottom  of  a  tall  vessel  and  fill  the 
vessel  with  a  liquid  in  which  it  is  soluble.  The  solid  will 
dissolve  and  pass  upward  through  the  solvent  to  the  surface. 

The  above  holds  true  for  the  heaviest  solids  and  the 
lightest  solvents.  The  substance  to  be  dissolved  may  be 
heavy,  such  as  a  salt  of  lead  or  platinum,  and  the  solvent 
may  be  water  or  some  much  lighter  liquid.  The  above  will 
take  place  regardless  of  the  relative  specific  gravities  of  the 
two.  This  is,  of  course,  the  well-known  phenomenon  of 
diffusion;  but  we  are  so  familiar  with  the  results  of  diffusion 
that  we  are  often  not  inclined  to  stop  and  think  of  its  cause. 

A  moment's  thought  will  show  that  this  is  a  very  remark- 
able phenomenon  —  a  heavy  salt  rising  against  the  pull 
of  gravity  through  a  very  much  lighter  liquid. 

Diffusion  Caused  by  Some  Force.  —  To  accomplish 
either  of  the  above  results  it  is  obvious  that  some  force  is 
necessary,  and  the  question  arises,  what  is  the  force  that 
causes  diffusion? 

If  we  examine  all  of  the  known  forces  in  solution,  we  will 
soon  come  to  the  conclusion  that  there  is  only  one  which 
can  possibly  cause  diffusion,  and  that  is  osmotic  pressure,      . 
or  the  force  which  causes  osmotic  pressure.    If  this  is  tfie 
only  known  force  in  solution  which  can  cause  diffusion,  why 


DIFFUSION  IN  SOLUTION  91 

not  conclude  at  once  that  the  cause  of  diffusion  is  osmotic 
pressure?  The  answer  is  of  course  obvious.  There  may  be 
in  solution  some  force  at  present  unknown  to  us  which  is 
the  cause  of  the  phenomenon  we  are  considering. 

Is  osmotic  pressure  the  cause  of  diffusion?  The  question 
can  be  answered  in  only  one  way.  Study  the  phenomena 
presented  by  osmotic  pressure  and  discover  their  laws. 
Study  the  phenomena  of  diffusion  and  discover  their  laws, 
and  then  see  whether  the  two  classes  of  phenomena  obey 
the  same  laws.  This  is  exactly  what  has  been  done  in  the 
cases  of  diffusion  and  osmotic  pressure.  With  what  result 
we  shall  now  see. 

Work  of  Graham  on  Diffusion.  —  The  pioneer  hi  the 
study  of  diffusion  hi  a  scientific  manner  was  the  English 
chemist,  Thomas  Graham.1  Graham's  method  of  studying 
diffusion  was  very  simple.  The  solution  was  poured  into  a 
glass  vessel,  and  water  poured  on  top  of  the  solution  until  the 
vessel  was  filled.  This  vessel  containing  the  solution  was 
then  placed  in  a  glass  cylinder,  which  was  filled  with  water 
until  the  level  of  the  water  in  the  cylinder  was  a  few  centi- 
meters above,  the  top  of  the  vessel  containing  the  solution. 
The  whole  system  was  then  set  in  a  quiet  place  at  a  constant 
temperature,  and  allowed  to  stand,  for  a  certain  definite 
period  of  time.  An  analysis  of  the  solution  hi  the  outer 
cylinder  gave  the  amount  of  the  dissolved  substance,  which, 
in  a  given  time,  under  the  given  conditions,  had  diffused 
from  the  inner  vessel  in^o  the  outer. 

Graham  found  that  the  mass  of  the  substance  that  dif- 
fuses is  largely  dependent  on  the  nature  of  the  substance. 
Acids  diffuse  much  morA-apidly  than  salts.  Salts,  hi  turn, 
diffuse  with  very  different  velocities.  With  mixtures  of 
salts  the  constituents  difi^sed  with  velocities  which  were 
characteristic  of  the  individual  salts;  and  each  diffused 
practically  as  if  it  alone  were  present. 

"If 2  two  salts  mix  without  combining  with  one  another, 

1  Phil  Trans.,  1,  805  (1850),  and  483  (1851).  Lieb.  Ann.,  77,  56  and  129 
(1851);  80,  197  (1851).  2  Lieb.  Ann.,  77,  75  (1851). 


92  THE  NATURE  OF  SOLUTION 

we  may  expect  that  both  will  diffuse  separately  and  inde- 
pendent of  one  another;  each  salt  obeying  its  own  specific 
power  of  diffusion." 

Graham  utilized  this  independent  diffusibility  of  salts 
from  their  mixtures  to  effect  separations  of  such  mixtures 
into  their  constituents.  "From1  what  has  been  said,  it  is 
clear  that  diffusion  furnishes  us  with  a  means  of  separating 
salts  to  a  certain  extent,  according  to  the  same  principle  by 
which  volatile  substances  of  different  boiling-points  can  be 
separated  by  distillation." 

This  applies  especially  to  salts  which  are  decomposed  by 
water,  such  as  acid  potassium  sulphate,  the  alums  and  other 
double  salts.  From  solutions  of  these  substances  the  con- 
stituents can  be  separated  to  a  greater  or  lesser  extent  by 
diffusion. 

Graham  also  studied  the  effect  on  the  diffusion  of  a  salt, 
of  dissolving  it  in  a  solution  of  another  salt.  He  wanted  to 
see  whether  the  diffusion  of  a  salt  in  a  solution  of  another 
salt  bore  any  relation  to  its  diffusion  in  pure  water.  He 
could  detect  no  difference  under  the  two  sets  of  conditions. 

While  Graham  showed  that  the  amount  of  diffusion  is 
approximately  proportional  to  the  concentration,  he  did  not 
arrive  at  the  fundamental  law  of  diffusion.  This  was  left 
for  Fick. 

The  impression  which  the  work  of  Graham  made  upon 
his  contemporaries  is  shown  by  the  following  editorial 
comment  on  the  first  paper  published  on  this  subject 
by  Graham.2 

"The  above  investigation  of  Professor  Graham  calls 
attention  to  the  existence  of  a  new  cause,  or  a  new  property 
of  heterogeneous  substances,  which  exerts  a  definite  effect 
on  the  affinity  hi  the  chemical  compound.  That  the  salt 
content  of  a  salt  solution  which  is  covered  with  pure  water, 
distributes  itself  gradually  from  below  towards  the  surface 
of  the  water,  is  a  well-known  fact,  but  that  different  sub- 
stances have  this  capability  to  a  very  unequal  degree,  that 

1  Ueb.  Ann.,  77,  76  (1851).  1  Ibid.,  77,  56  (1851). 


DIFFUSION  IN  SOLUTION  93 

each  obeys  its  own  specific  law  of  diffusion,  that  in  conse- 
quence of  this  inequality  the  water  supernatant  to  a  solu- 
tion of  mixed  salts  receives  the  mixed  salts  in  a  different 
proportion  from  that  hi  which  they  exist  in  the  lower  layers, 
that  under  these  conditions,  alum,  hi  consequence  of  the 
unequal  diffusibility  of  its  constituents,  is  decomposed, 
that  due  to  the  same  cause  a  solution  of  potassium  sulphate 
in  lime  water  is  decomposed;  all  these  are  such  remarkable 
and  unexpected  results  that  they  are  to  be  ranked  among 
the  most  important  contributions  to  science  hi  recent  tunes." 
This  note  is  signed  "The  editor";  and  the  "Annalen"  was 
edited  at  that  tune  by  Friedrich  Wohler,  Justus  Liebig  and 
Hermann  Kopp. 

The  Generalization  of  Fick.  —  Fick  studied  the  diffusion 
of  solutions  with  special  reference  to  the  relation  between 
the  amounts  of  substances  that  diffuse,  the  concentra- 
tions of  the  solutions  and  the  differences  in  concentrations 
between  diffusing  solutions.  The  result  is  the  generalization 
known  as  the  law  of  Fick.1  This  law  may  be  formulated 
as  follows: 

"The2  dissemination  of  a  dissolved  substance  through  a 
solvent,  as  far  as  it  takes  place  undisturbed  under  the 
influence  of  molecular  forces  alone,  obeys  the  same  laws 
which  Fourier  established  for  the  dissemination  of  heat  hi  a 
conductor,  and  which  Ohm  announced  for  the  conduction  of 
electricity.  In  the  Fourier  law  it  is  only  necessary  to  replace 
the  words  quantity  of  heat  with  the  words  quantity  of 
dissolved  substance,  and  the  word  temperature  with  density 
of  solution." 

In  a  word,  Fick  showed  that  diffusion  is  proportional  I 
to  difference  hi  concentration. 

Effect  of  Mass  on  Diffusion.  —  From  the  law  of  Fick 
it  can  be  seen  at  once  that  the  result  of  diffusion  is  to  make 
the  whole  system  homogeneous.  No  matter  what  the  origi- 
nal difference  in  concentration  may  have  been,  no  matter 
what  the  nature  of  the  substance  or  substances  diffusing,  no 

1  Pogg.  Ann.,  94,  59  (1855).  2  Ibid.,  94,  65  (1865). 


94  THE  NATURE  OF  SOLUTION 

matter  what  the  nature  of  the  solvent  or  solvents  involved; 
the  dissolved  substance  diffuses  over  from  the  region  of 
greater  to  that  of  lesser  concentration,  and  this  continues 
until  equality  of  concentration  is  established  hi  all  parts  of 
the  solution.  This  must  be  regarded  as  one  of  the  funda- 
mental facts  in  connection  with  solution.  If  this  fact  is 
considered  with  that  to  which  attention  was  called  earlier  in 
this  paper,  that  diffusion  is  a  phenomenon  which  seems 
largely  to  defy  gravitation,  we  realize  the  remarkable  nature 
of  the  property  with  which  we  are  dealing. 

When  we  think  of  very  heavy  salts  rising  vertically 
through  comparatively  light  liquids,  however,  we  cannot 
avoid  the  question  as  to  whether  the  mass  of  the  diffusing 
substance  does  not  have  something  to  do  with  its  diffusibility, 
if  not  with  the  final  conditions  of  equilibrium  of  the  diffus- 
ing system. 

Beudant,1  in  1818,  announced  that  he  had  made  the 
observation  that  when  unsaturated  solutions  were  allowed 
to  stand,  crystals  separated  on  the  bottoms  of  vessels  con- 
taining such  solutions.  This  observation,  if  correct,  would 
be  of  the  very  greatest  importance  as  bearing  on  the  problem 
in  hand. 

It  attracted  the  attention  of  Gay-Lussac,2  who  at  once 
repeated  the  experiment.  He  quotes  Beudant,3  "I  have 
remarked  that  crystals  are  formed  without  any  evaporation 
from  solutions  otherwise  very  dilute."  Beudant  states  that 
he  took  proper  precautions  to  prevent  any  evaporation  of 
the  solutions. 

Says  Gay-Lussac,4  "I  have  taken  two  glass  tubes  two 
meters  long  and  three  centimeters  in  diameter.  I  placed 
hi  -one  a  saturated  solution,  of  niter,  at  the  temperature  of 
the  vault  of  the  observatory,  and  in  the  other  a  solution 
of  marine  salt,  also  saturated;  two  other  tubes  were  filled 
with  similar  solutions  in  which  there  were  not  more  than 

1  Ann.  Chim.  Phys.  [2],  8,  15  (1818). 

2  Ibid.  [2],  11,  306  (1819). 

8  Ibid.  [2],  8,  15  (1818).  <  Loc.  tit. 


DIFFUSION  IN  SOLUTION  95 

four  centigrams  of  each  salt.  These  tubes,  hermetically 
sealed,  remained  six  months  hi  the  vaults  of  the  observatory 
hi  a  vertical  position.  At  the  end  of  this  tune  I  determined 
by  evaporation  the  quantity  of  salt  contained  hi  the  water 
of  the  upper  and  of  the  lower  part  of  the  tube,  and  found 
that  the  solutions  were  perfectly  homogeneous." 

Gay-Lussac  then  attempts  to  explain  the  source  of  error 
hi  Beudant's  experiments.  It  would  lead  us  too  far  to 
discuss  this  here,  and  it  is  not  necessary,  since  they  are 
probably  wrong. 

In  a  paper  by  Lieben,1  "On  the  Homogeneity  of  Solu- 
tions," he  studies  the  above  question  for  solutions  of  both 
solids  and  gases.  He  states  that  his  results  justify  the  con- 
clusion that,  "So2  far  as  our  present  knowledge  goes,  hi 
a  homogeneous  solution  which  is  allowed  to  rest  quietly, 
a  settling  of  the  particles  of  salt  from  the  upper  to  the 
lower  layers  does  not  take  place." 

This  same  question  has  been  discussed  thermodynamically 
by  Gouy  and  Chaperon,  and  the  conclusions  tested  by 
experimental  results;  hi  a  paper  on  the  concentrating  of 
solutions  by  weight,3  they  show  that  weight  has  no  effect 
provided  the  solution  does  not  change  hi  density  when  sub- 
jected to  small  changes  hi  concentration.  If  the  density 
increases  with  increase  hi  concentration,  the  lower  layers 
become  more  concentrated. 

From  the  vapor-tension  measurements  of  Wullner  and 
Moser4  the  authors  calculate  the  difference  hi  concentra- 
tion between  the  uppermost  and  lowest  layers  of  a  solution 
100  meters  deep. 

They  obtained  the  following  results: 

At  top  At  bottom  Difference 

Cadmium  iodide  0.166  0.153  0.013 

Sodium  nitrate  0.20  0.196  0.004 

Sodium  chloride  0.11  0.1095  0.0005 

Sugar  0.55  0.546  0.004 

1  Lieb.  Ann,,  101,  77  (1857).  *  Ibid.,  83. 

3  Ann.  Chim.  Phys.  [6],  12,  384  (1887). 

4  Pogg.  Ann.,  103,  Wied.  Ann.,  3,  (1881). 


96  THE  NATURE  OF  SOLUTION 

These  differences  are  obviously  very  small. 

Temperature  Coefficients  of  Diffusion  —  Principle  of 
Soret.  —  Diffusion  as  it  takes  place  ordinarily  is  a  very  slow 
process.  It  requires  months  and,  under  certain  conditions, 
many  months  to  establish  equilibrium  where  diffusion  alone 
establishes  the  equilibrium.  This  condition  is  accelerated 
by  rise  in  temperature  —  the  higher  the  temperature,  other 
things  being  equal,  the  more  rapid  the  diffusion.  What 
are  the  temperature  coefficients  of  diffusion? 

If  we  take  a  homogeneous  solution  and  keep  the  dif- 
ferent parts  at  the  same  temperature,  the  solution  will 
remain  homogeneous.  If,  on  the  other  hand,  the  different 
parts  of  a  homogeneous  solution  are  kept  at  different  tem- 
peratures, the  solution  will  no  longer  remain  homogeneous, 
but  the  different  parts  will  have  different  concentrations  — 
the  colder  parts  of  the  solution  will  become  more  concen- 
trated and  the  warmer  parts  more  dilute.  This  was  first 
observed  by  Ludwig1  in  1856: 

This  phenomenon  was  thoroughly  studied  by  Soret,2  as 
has  been  discussed  under  the  " Principle  of  Soret"  (p.  75). 
The  change  in  concentration,  as  we  saw,  when  equilibrium 
is  established  is  calculable  from  the  law  of  Gay-Lussac, 
and  the  difference  in  concentration  between  the  colder 
and  the  warmer  parts  of  the  solution  is  expressed  thus. 
Suppose  the  colder  part  of  the  solution  is  at  temperature  a, 
and  the  warmer  part  at  temperature  6,  the  difference  in 

concentration  is  -^=^- 
27o 

Osmotic  Pressure  Produces  Diffusion.  —  We  now  come 
to  the  fundamental  point  in  connection  with  diffusion. 
What  causes  it? 

A  dissolved  substance  passes  from  one  part  of  a  solution 
to  another.  The  different  parts  of  the  solution  may  be 
widely  removed  from  one  another,  one  part  may  be  very 
high  above  the  other;  and  yet  the  dissolved  substance  will 

i  Wien.  Ber.  20,  539  (1856). 

8  Ann.  Chim.  Phys.  [5],  22,  293  (1881). 


DIFFUSION  IN  SOLUTION  97 

rise  to  the  top  of  the  solvent,  as  we  have  seen,  right  against 
the  pull  of  gravitation. 

It  is  obvious  that  to  effect  such  results  some  force  is 
necessary,  and  it  must  be  a  force  of  very  considerable  mag- 
nitude. The  question  is,  what  is  this  force? 

On  thinking  over  the  forces  which  exist  hi  solution  the 
most  obvious,  as  we  have  seen  (p.  90),  which  might  pro- 
duce diffusion  is  osmotic  pressure  —  does  it? 

This  question  can  be  answered  most  directly  by  com- 
paring the  laws  of  osmotic  pressure  with  the  laws  of  diffusion. 
A  moment's  thought  will  show  that  the  law  of  Fick  for  dif- 
fusion is  strictly  analogous  to  the  law  of  Boyle  for  osmotic 
pressure.    Fick's  law  says  that  the  amount  of  substance  j 
that  will  diffuse  from  the  one  solution  to  the  other  is  pro-  J 
portional  to  the  difference  hi  concentration  between  the  , 
two  solutions.    Boyle's  law  for  osmotic  pressure  states  that 
the  osmotic  pressure  at  the  surface  of  contact  of  two  solu- 
tions, or  of  a  solution  and  a  solvent,  is  proportional  to  the 
difference  hi  concentration  of  the  two  solutions.    The  two 
phenomena  thus  obey  exactly  the  same  law  as  far  as  the 
quantity  of  substance  that  diffuses  and  the  amount  of  os- 
motic pressure  under  the  same  conditions  are  concerned. 

Having  found  that  the  law  of  Fick  for  diffusion  is  analo- 
gous to  the  la,w  of  Boyle  for  osmotic  pressure,  we  would 
naturally  look  farther  and  see  whether  there  are  any  other 
laws  of  diffusion  which  can  be  compared  with  laws  of  osmotic 
pressure.  The  most  obvious  for  such  further  comparisons  is 
the  law  of  the  temperature  coefficients  of  the  two  phenomena. 

It  will  be  recalled  that  the  law  of  Gay-Lussac  for  gases 
was  shown,  from  the  measurements  of  osmotic  pressure  by 
Pfeffer,  to  hold  approximately  for  solutions.  The  recent 
measurements  of  osmotic  pressure  by  Morse  and  his  co- 
workers  have  shown  that  this  law  holds  approximately  for 
osmotic  pressure  under  all  conditions,  and  rigidly  under 
many  conditions.  The  principle  of  Soret  for  diffusion 
snows  that  the  law  of  Gay-Lussac  holds  at  least  approxi- 
mately for  the  temperature  coefficients  of  diffusion;  and  the 


98  THE  NATURE  OF  SOLUTION 

most  recent  work,  in  which  the  solutions  were  allowed  to 
stand  for  a  long  time,  indicates  that  this  law  holds  very 
closely  or  perhaps  rigidly  for  diffusion.  Two  of  the  funda- 
mental laws  of  osmotic  pressure  therefore  apply  to  the 
phenomena  of  diffusion,  and  we  would  naturally  conclude 
either  that  the  one  is  the  cause  of  the  other,  or  that  diffu- 
sion and  osmotic  pressure  are  both  the  results  of  a  common 
cause,  which  at  present  is  entirely  unknown.  But  since  we 
know  of  no  such  common  cause  we  may  say  that  the  cause 
of  diffusion  is  osmotic  pressure.  1 

Let  us  recall  the  order  of  magnitude  of  osmotic  pressure 
(p.  54),  and  we  see  that  we  are  dealing  with  a  very  large 
force.  A  normal  aqueous  solution  of  a  nonelectrolyte  hi 
contact  with  pure  water  exerts  an  osmotic  pressure  of 
between  twenty  and  thirty  atmospheres.  Many  sub- 
stances are  so  soluble  that  we  can  prepare  solutions  five  or 
even  ten  times  normal.  Such  solutions  would,  respectively, 
have  at  least  one  hundred  and  twenty-five  and  two 
hundred  and  fifty  atmospheres  of  osmotic  pressure,  when 
brought  in  contact  with  the  pure  solvent.  If  the  dissolved 
substances  were  electrolytes  and  therefore  dissociated, 
their  solutions  would  show  still  greater  osmotic  pressures. 
If  the  solute  were  a  binary  electrolyte  each  molecule  would 
dissociate  into  two  ions;  if  a  ternary  electrolyte,  into  three 
ions;  and  so  on.  Since,  as  we  have  seen,  an  ion  exerts  the 
same  osmotic  pressure  as  a  molecule,  such  solutions,  if  com- 
pletely dissociated,  would  exert  two  or  three  times  the 
osmotic  pressures  of  solutions  of  nonelectrolytes  of  corre- 
sponding concentrations. 

We  can  gain  from  these  facts  some  idea  of  the  magnitude 
of  the  force  which  we  call  osmotic  pressure.  For  solutions 
of  any  appreciable  concentration,  it  is  large;  for  concen- 
trated solutions,  it  is  enormous.  This  shows  us  why  it  is 
that  dissolved  substances  diffuse  practically  independent  of 
the  relative  specific  gravities  of  solvent  and  solute.  The 
driving  force  is  so  great  that  the  relative  densities  of  the 
dissolved  particle  and  solvent  are  small  in  comparison. 


DIFFUSION  IN  SOLUTION  99 

Importance  of  Osmotic  Pressure  for  Chemistry.  —  Now 
that  we  have  seen  that  osmotic  pressure  is  the  probable 
cause  of  diffusion,  we  can  begin  to  realize  the  importance 
and  significance  of  osmotic  pressure  for  nature,  and  therefore 
for  the  science  of  nature  or  natural  science. 

If  osmotic  pressure  did  not  exist  we  would  not  have 
diffusion  hi  solution.  If  we  did  not  have  diffusion,  we  could 
not  have  a  homogeneous  solution  for  any  appreciable  length 
of  time,  unless  we  kept  it  constantly  hi  a  violent  state  of 
agitation.  It  would  therefore  be  impossible  to  maintain  a 
solution  standard  for  any  time.  Without  standard  solu- 
tions all  volumetric  analysis  would,  of  course,  be  impossible. 
Without  volumetric  analysis  much  of  chemistry  would  be 
changed,  since  there  are  many  analytical  processes  that  we 
can  carry  out  volumetrically,  but  we  cannot  gravimetrically. 
Remove  standard  solutions  from  chemistry  and  this  branch 
of  science  would  be  set  back  indefinitely. 

Bearing  of  Osmotic  Pressure  on  Biological  Phenomena. 
—  When  we  turn  to  biology,  we  find  that  homogeneous 
solutions  are  almost  as  important  as  for  chemistry.  Living 
matter  must  often  be  surrounded  by  a  solution  of  a  certain 
substance  or  of  certain  substances,  having  a  concentration 
which  is  just  equal  to  and  in  no  case  exceeds  a  certain  value. 
If  osmotic  pressure,  and  consequently  diffusion,  did  not  exist, 
such  solutions  on  standing  would  become  more  concentrated 
hi  some  parts  and  more  dilute  in  others.  Those  living 
forms  which  depend  for  their  existence  upon  a  definite 
concentration  of  certain  substances,  would  cease  to  live 
as  soon  as  the  homogeneity  of  the  solution  was  destroyed, 
one  part  becoming  too  concentrated  for  their  existence, 
the  other  too  dilute. 

Approximate  homogeneity  of  solution  is  often  of  funda- 
mental importance  in  obtaining  food.  This  applies  especially 
to  nonmobile  animals  and  to  plants  which  cannot  go  after 
then*  food,  but  which  must  depend  upon  the  food  being 
brought  to  them  in  solution.  If  such  solutions  were  non- 
homogeneous,  being  more  concentrated  in  certain  parts 


100  THE  NATURE  OF  SOLUTION 

than  in  others,  there  would  be  an  excess  of  food  in  certain 
parts  and  little  or  none  in  other  portions  of  the  solution. 
The  result  is  obvious.  Many  individuals,  being  deprived 
of  sufficient  food,  would  die. 

If  there  were  no  such  force  in  a  solution  to  produce  or 
maintain  homogeneity,  those  parts  of  a  solution  of  food- 
stuffs which  under  one  set  of  conditions  were  the  more 
concentrated,  under  other  conditions  might  contain  little  or 
no  food.  Those  individuals  which  at  one  time  might  be 
most  favorably  placed  with  respect  to  food,  at  another  time 
might  be  completely  deprived  of  it,  and,  consequently,  ex- 
terminated. In  this  way  it  is  quite  easy  to  conceive  of  a 
whole  species  being  exterminated,  due  to  the  lack  of  a  fairly 
homogeneous  solution  of  food. 

Then  consider  an  even  more  fundamental  question  for 
living  forms  than  that  of  food  supply;  i.e.,  the  contents 
of  the  cells  themselves  of  which  all  living  forms  are  made. 
These  cells  are  often  very  sensitive  systems.  A  slight 
disturbance  either  in  the  solutions  which  surround  them 
or  still  more,  perhaps,  in  their  own  contents,  would  produce 
death.  Let  us  suppose  that  the  fluids  surrounding  the 
cells,  which  are  usually  solutions  of  a  great  number  of 
things,  should  no  longer  preserve  homogeneity  of  concen- 
tration, but  should  become  in  some  parts  more  concentrated 
and  hi  others  more  dilute,  then  what  would  happen?  The 
cell  could  live  and  function  normally  only  when  sur- 
rounded by  solutions  of  certain  substances  at  certain  definite 
concentrations.  The  solution  around  the  cell  has  now 
become  in  one  part  more  concentrated,  and  in  the  other  more 
dilute,  than  the  normal.  The  result  is  obvious.  The  cells 
in  the  more  dilute  part  of  the  solution  would  die  because 
the  solution  was  too  dilute  for  their  existence.  Those  in 
the  more  concentrated  part  of  the  solution  would  die  be- 
cause the  solution  was  too  concentrated  for  them  to  per- 
form the  normal  functions  of  life.  The  result  would  be 
disastrous  for  all  the  cells. 

Take  one  more  step  and  go  within  the  cell  itself.    A  cell 


DIFFUSION  IN  SOiAjfHMT' 

is,  for  our  purpose,  an  aqueous  solution  of  a  large  number 
of  substances  each  usually  having  a  pretty  definite  con- 
centration: the  whole  surrounded  by  a  sac.  The  cell  is, 
further,  a  very  sensitive  system.  Slight  disturbances  in  the 
normal  equilibrium  of  concentrations  of  the  various  constit- 
uents will  frequently  produce  death. 

Suppose  the  solutions  within  the  cell  did  not  remain 
homogeneous,  but  that  the  heavier  constituents  settled  to 
the  bottom,  as  would  be  the  case  if  there  were  no  diffusion  to 
maintain  homogeneity  and  no  osmotic  pressure  to  produce 
diffusion.  The  result  is  obvious.  The  cell  would  die. 

We  thus  see  that  homogeneity  of  solutions,  both  without 
and  within  the  cells,  is  absolutely  essential  to  living  matter. 
This  homogeneity  can,  of  course,  be  maintained  only  by  dif- 
fusion, and  diffusion  is  due  to  osmotic  pressure.  We  thus 
see  that  osmotic  pressure  hi  solutions  is  absolutely  essential 
to  life  as  we  now  know  it.  Indeed,  it  is  difficult  to  conceive 
of  living  matter  having  any  close  relation  to  life  as  it  now 
exists  upon  the  earth,  without  osmotic  pressure  hi  solution 
to  maintain  homogeneity  of  concentration. 

Osmotic  Pressure  and  Geology.  —  When  we  turn  from 
biology  to  geology  the  effect  of  osmotic  pressure  hi  maintain- 
ing solutions  of  homogeneous  concentration  is  hardly  less 
important.  Take,  for  example,  the  rocks  that  are  deposited 
from  aqueous  solutions.  If  such  solutions  were  not  main- 
tained more  or  less  homogeneous  by  diffusion,  the  compo- 
sitions of  the  various  deposits  from  such  solutions  would 
be  very  different  from  what  we  now  know  them  to  be. 
The  heavier  constituents,  being  near  the  bottoms  of  the  solu- 
tions, would  tend  to  come  down  first,  and  the  present  order 
of  such  deposits  would  be  greatly  changed. 

In  the  case  of  the  igneous  rocks  the  effects  of  hetero- 
geneous solutions  would  be  perhaps  even  more  pronounced. 
The  molten  magmas,  many  of  which,  as  we  have  seen,  are 
true  solutions,  would  deposit  rocks  of  very  different  compo- 
sition, if  there  were  no  approach  to  homogeneity  in  the  com- 
position of  such  magmas.  If  a  magma,  once  approximately 


102  TtiE  NATURE  OF  SOLUTION 

homogeneous,  were  to  become  very  heterogeneous  due  to  the 
absence  of  all  diffusion,  the  heavier  constituents  of  such  a 
magma  would,  of  course,  tend  to  settle  to  the  bottom  of  the 
fused  mass,  and  when  this  had  become  sufficiently  cool, 
would  solidify  and  form  rocks  of  very  different  composition 
from  those  now  known  to  us  as  having  been  deposited  from 
such  magmas. 

Without  diffusion  and  its  cause,  osmotic  pressure,  the 
composition  of  the  rocks  would  be  very  different  from  what 
it  is,  and  the  rocks  are  the  elements  of  geology. 


CHAPTER  VII 

DEPRESSION    OF    THE    VAPOR-TENSION    OF   A    SOLVENT     BY 
SUBSTANCES   DISSOLVED    IN   IT 

IT  has  long  been  known  that,  hi  general,  when  solids 
are  dissolved  in  liquids,  the  vapor-tensions  of  the  latter 
at  any  given  temperature  are  reduced.  When  liquids  are 
dissolved  hi  liquids  the  same  phenomenon  manifests  itself, 
provided  the  boiling-point  of  the  liquid  solute  is  consid- 
erably higher  than  that  of  the  solvent. 

Work  of  Faraday.  —  The  first  to  have  investigated  this 
phenomenon  quantitatively  was  Michael  Faraday,  all 
things  considered,  one  of  the  very  greatest  experimenters 
who  has  ever  lived.  In  1822  he  published  a  paper1  on  the 
"  Temperature  Produced  by  the  Condensation  of  Vapor." 
He  determined  the  boiling-points  of  saturated  solutions  of 
a  number  of  salts:  acid  potassium  carbonate,  potassium 
tartrate,  ammonium  chloride,  etc.  From  his  work  Faraday2 
concluded  that  "the  temperature  of  aqueous  vapor  from  a 
saline  solution  is  always  100°,  whatever  the  boiling-point 
of  the  solution  under  atmospheric  pressure." 

The  editor,  Gay-Lussac,3  makes  the  following  comment, 
which,  when  we  consider  the  men  involved,  has  a  particular 
interest. 

"It  would  be  very  difficult  to  conceive  of  the  vapor 
given  off  from  the  surface  of  a  salt  solution,  not  having 
exactly  the  same  temperature  as  the  solution  from  which  it 
came.  Without  introducing  any  theory,  we  can  state  from 
the  irrefutable  evidence  of  experiment  that  the  temperature 
of  the  vapor  from  any  liquid  whatever,  under  any  pressure 
whatever,  is  exactly  equal  to  that  of  the  layer  of  liquid  imme- 

1  Ann.  Chim.  Phys.  [2],  20,  320  (1822). 

2  Ibid.,  325.  3  Ibid.,  325. 


104  THE  NATURE  OF  SOLUTION 

diately  in  contact  with  the  vapor.  If  this  fact  escaped  Faraday, 
we  are  convinced  that  it  is  because  he  had  not  made  a 
sufficiently  large  number  of  experiments,  and  had  not  worked 
with  a  sufficiently  sensitive  thermometer." 

Investigations  of  Wiillner.  —  The  next  work  in  this  field 
of  special  interest  and  significance  is  that  of  Wiillner,1 
which  was  done  between  three  and  four  decades  after  that 
of  Faraday. 

A  few  of  the  results  obtained  by  Wiillner  will  serve  to 
illustrate  the  relation  discovered  by  him.  If  we  represent 
the  vapor-pressure  of  pure  water  by  P,  the  lowering  of  the 
vapor-pressure  of  100  parts  of  water  by  one  part  of  salt  by 
PI,  we  have  the  following  relations.  Wiillner's  results  for 
only  one  salt,  sodium  chloride,  will  be  given. 

P  5%  10%  20%  30%  P!             _Pi 

Temperature      mm.  mm.  mm.  mm.  mm.        Mean          P 

19.9°              17.28  0.98  1.47  4.06       0.149 

35.0°              41.82  1.79  3.11  8.13       0.300  0.0072 

49.8°              91.09  2.84  5.81  11.36  18.20       0.620  0.0068 

64.8°  185.27  —  9.88  21.45  1.090  0.0059 

82.2°  388.33  —  21.58  46.20  2.240  0.0058 

100.5°  775.40  —  44.90  92.41  4.450  0.0057 

p 

We  see  from  the  above  table  that  -^  is  practically  con- 
stant, if  we  consider  the  comparatively  large  experimental 
errors  involved  in  the  results.  From  his  study  of  a  fairly 
large  number  of  salts  Wiillner  generalized  his  results  as 
follows: 

"The  lowering  of  the  vapor-pressure  of  water  by  dis- 
solved substances  is  proportional  to  the  amounts  of  sub- 
stances in  the  solution." 

This  conclusion  was  called  in  question  about  twenty 
years  later  by  Pauchon,2  and  Tammann3  confirmed  the 
validity  of  Pauchon's  objection.  The  latter  showed  that 
the  so-called  law  of  Wiillner  was  only  an  approximation. 

Study  of  Vapor-Tension  by  Walker.  —  The  next  really 

1  Pogg.  Ann.,  103,  529  (1858);  105,  85  (1858);  110,  564  (I860). 

2  C&mpt.  Rend.,  89,  752  (1879). 

3  Wied.  Ann.,  24,  523  (1885). 


DEPRESSION  OF  THE  VAPOR-TENSION  105 

important  investigation  in  this  field  is  that  of  Walker,1 
carried  out  in  1888  hi  the  laboratory  of  Ostwald  in  Leipzig. 

His  method  of  work  was  probably  the  best  that  had  been 
used  up  to  that  time,  and  is  interesting.  A  few  words 
concerning  it  are  worth  while. 

The  problem  was  to  compare  directly  the  vapor-tension 
of  water  with  that  of  an  aqueous  solution.  The  apparatus 
was  very  simple,  consisting  of  three  ordinary  Liebig  bulbs, 
such  as  are  used  in  the  combustion  analysis  of  organic 
compounds  and  a  U-tube.  The  solution  was  introduced 
into  two  of  the  bulbs,  while  the  third  contained  the  pure 
solvent.  Fragments  of  pumice  stone  moistened  with  sul- 
phuric acid  were  added  to  the  U-tube. 

Air  dried  by  sulphuric  acid  was  drawn  slowly  through 
the  whole  system.  In  passing  through  bulbs  one  and  two 
the  air  was  saturated  with  water- vapor  at  the  tension  of  the 
vapor  over  the  solution.  When  this  air  was  drawn  through 
bulb  three  it  took  up  more  water-vapor,  since  the  tension 
of  the  vapor  of  water  was  greater  over  the  pure  solvent 
than  over  the  solution.  The  air  was  then  passed  through 
the  U-tube  containing  sulphuric  acid  and  dried  to  the 
same  degree  as  when  it  entered  the  first  tube  filled  with 
the  solution. 

The  loss  hi  weight  of  the  two  tubes  containing  the  solu- 
tion was  then  determined.  The  loss  hi  weight  of  the  tube 
containing  the  pure  solvent  was  also  determined.  The  gain 
hi  weight  of  the  tube  containing  the  sulphuric  acid  must  be 
exactly  equal  to  the  total  losses  in  the  weights  of  the  other 
three  tubes.  From  the  amount  of  water  lost  by  the  solu- 
tion, and  the  amount  lost  by  the  pure  solvent,  we  can  cal- 
culate directly  the  lowering  of  the  vapor-tension  of  the 
solvent  produced  by  the  dissolved  substance. 

A  few  of  the  results  obtained  by  Walker2  will  be  given. 
The  number  of  parts  salt  to  100  parts  water  hi  the  column 
under  g ;  c  is  the  relative  lowering  of  vapor-pressure  produced 
by  one  molecule  of  substance  in  100  molecules  of  water. 

1  Zeti.  phys.  Chem.,  2,  602  (1888).       *  Ibid.,  2,  604  (1888). 


106  THE  NATURE  OF  SOLUTION 

NaCl  5.96  2.07 

NaCl  18.60  2.18 

NaCl  32.265  2.29 

KC1  7.66  2.02 

NaNOa  8.791  1.97 

CaCl2  11.386  3.30 

MgCl2  4.791  3.57 

MgCl2  9.691  4.01 

BaCl2  21.443  3.00 

SrCl,  16.002  3.28 

One  conclusion  is  justified  from  these  results.  Where 
more  than  one  concentration  of  a  substance  was  studied, 
the  more  concentrated  the  solution  the  greater  the  relative 
lowering  of  the  vapor-tension.  This  fact  is  of  importance 
in  connection  with  a  theory  of  solution  which  will  be  de- 
veloped in  the  last  chapters  of  this  book. 

This  brings  us  to  the  most  important  series  of  investiga- 
tions that  has  thus  far  been  carried  out  on  the  vapor-tensions 
of  solutions  and  solvents;  viz.,  the  work  of  Raoult. 

Raoult  and  the  Vapor-Tensions  of  Solvents  and  of 
Solutions.  —  The  earlier  workers  in  this  field  had  studied 
aqueous  solutions  almost  exclusively.  This  we  know  today 
was  unfortunate,  and  for  two  independent  reasons.  In  the 
first  place,  water  at  ordinary  temperatures  has  compara- 
tively small  vapor-tension,  and  on  account  of  its  small 
molecular  weight  the  depression  of  its  vapor-tension  by  dis- 
solved substances  is  very  small.  We  are  then  measuring  a 
small  change  in  a  small  quantity,  and  the  experimental  error 
is,  of  course,  relatively  large. 

Further,  as  we  have  seen,  water  is  one  of  the  best  dis- 
sociating solvents.  This  means  that  electrolytes  dissolved 
in  this  solvent  are  largely  broken  down  into  ions.  Any 
relations  between  the  lowering  of  the  vapor-tension  and  the 
concentration  of  the  solution  which  might  exist,  would  in 
aqueous  solution  be  liable  to  be  masked  by  the  breaking 
down  of  the  molecules  into  ions  by  this  solvent. 

Raoult1  fortunately  turned  to  nonaqueous  solutions  and 
the  result  was  the  discovery  of  certain  relations  of  funda- 
mental significance.  Ether  is  a  solvent  which  has  a  high 

1  Ann.  Chim.  Phys.  [6],  15,  375  (1888). 


DEPRESSION  OF  THE  VAPOR-TENSION  107 

vapor-tension  even  at  ordinary  temperatures.  With  this 
solvent  the  quantity  to  be  measured  is  relatively  large,  and 
the  error  of  measurement  therefore  relatively  small.  Fur- 
ther since  ether  has  a  comparatively  large  molecular  weight, 
the  lowering  of  its  vapor-tension  by  dissolved  substances  is 
relatively  great,  as  we  shall  see.  It  was  therefore  an  ideal 
solvent  with  which  to  study  vapor-tensions  of  both  sol- 
vent and  solutions. 

The  experimental  work  of  Raoult  is  admirable  for  its 
insight  into  sources  of  experimental  error  and  ingenuity  in 
overcoming  or  correcting  for  these  errors.  The  method 
employed  was  the  barometric.  The  solvent  or  solution 
was  introduced  into  a  glass  tube  over  mercury  and  the  pres- 
sure determined  from  the  depression  of  the  mercury  column. 
The  observed  depression  was  corrected  for  a  number  of 
conditions  —  for  the  pressure  of  the  liquid  on  the  column 
of  mercury  beneath  it;  for  the  concentrating  of  the  solution 
at  and  near  its  surface,  due  to  the  loss  of  the  solvent  as 
vapor;  and  for  a  number  of  other  errors,  but  for  these, 
reference  must  be  had  to  the  original  paper.  Let  us  now 
see  what  Raoult  found. 

Raoult's  Results  —  Effect  of  Concentration.  —  Raoult 
took  up  first  the  effect  of  concentration  of  the  solution  on  the 
vapor-pressure  of  solutions  hi  ether.  If  we  represent  by  0 
the  number  of  grams  of  solute  in  100  grams  of  solvent;  by 
N,  the  number  of  molecules  of  solute  in  100  molecules  of 
solvent,  readily  calculated  from  0,  knowing  the  molecular 
weights  of  solute  and  solvent;  by  R,  the  ratio  between  the 
molecular  weights  of  solute  and  solvent  multiplied  by  100; 
we  have,  for  solutions  of  oil  of  turpentine  in  ether  the  fol- 
lowing results.1 

R  R 

0                             N  Observed  Calculated 

10.2                             5.9                                94.0  94.7 

20.2  12.1                                88.1  89.1 
35.9  23.4                                78.1  78.9 

50.3  35.5                                67.6  68.0 

62.8  47.9  56.2  56.9 

76.9  64.5  42.1  42.0 

1  Taken  from  Ann.  Chim.  Phys.  [6],  15,  388  (1888). 


108  THE  NATURE  OF  SOLUTION 

The  values  in  the  last  column  are  calculated  from  the 
equation  /'  74  X  p' 

f~  ^74Xp' 


in  which  /'  is  the  vapor-pressure  of  the  solution,  /,  that  of 
pure  ether,  p',  the  weight  of  the  dissolved  substance,  p,  the 
weight  of  the  ether,  m,  the  molecular  weight  of  the  dissolved 
substance,  and  74  the  molecular  weight  of  ether;  K  a  co- 
efficient whose  value  depends  upon  the  nature  of  the  solute. 

The  agreement  between  the  observed  and  calculated 
values  is  surprisingly  good.  Equally  satisfactory  agree- 
ments between  observed  and  calculated  values  were  found 
for  nitrobenzene  in  ether,  aniline  in  ether,  etc. 

The  value  of  the  coefficient  K  for  various  substances 
dissolved  in  ether  is  very  nearly  constant,  as  the  following 
results  will  show. 

Oil  of  turpentine  K  -  0.90 

Aniline  K  -  0.90 

Ethyl  benzoate  K  =  0.90 

Methyl  salicylate  K  =  0.82 

Nitrobenzene  K  =  0.70 

For  other  solvents  K  is  only  a  little  less  than  unity. 

Having  formulated  the  effect  of  concentration  on  the 
lowering  of  the  vapor-tension  of  solvents  by  dissolved  sub- 
stances, Raoult  raised  and  answered  the  question  as  to  the 
effect  of  temperature  on  the  relative  lowering  of  the  vapor- 
tension  of  ethereal  solutions. 

Effect  of  Temperature.  —  Raoult  dissolved  four  high- 
boiling  substances  hi  ether,  and  measured  carefully  the 
vapor-pressures  above  these  solutions  at  different  tempera- 
tures. The  range  of  temperature  was  from  0°  to  21°.  The 
following  results  were  obtained  by  him,  /  being  the  vapor- 
tension  of  the  solvent,  and/'  that  of  the  solution.1 

27.601  GRAMS  HEXACHLORETHANE  IN  100  GRAMS  ETHER 
Temperature  /  /'  £  X  100 

1.0°  197.0  181.3  92.0 

3.7°  224.2  205.4  91.6 

18.8°  418.6  280.9  91.0 

21.0°  457.3  417.8  91.4 

1  Taken  from  Ann.  Chim.  Phys.  [6],  15,  388  (1888). 


DEPRESSION  OF  THE  VAPOR-TENSION  109 

f 

The  variation  of  —  from  a  constant  is  scarcely  greater 

than  the  experimental  error,  over  the  temperature  range 
0°  to  21°. 

Effect  of  Nature  of  Dissolved  Substance.  —  The  most 
important  question  still  remains  —  the  effect  of  the  nature 
of  the  dissolved  substance  on  the  lowering  of  the  vapor-tension 
of  ether.  Raoult  deduced  the  following  very  simple  and 
important  relation:  if  we  represent  the  vapor-tension  of 
ether  by  /,  and  of  ethereal  solutions  by  /',  the  number  of 
molecules  of  solute  in  100  molecules  of  solution  by  N,  we 
have, 

j  X  100  =  100  -  KN 
f-f     KN 

or>  "T"  =Ioo 

The  quantity  —7—  is  called  by  Raoult  the  "  relative  dimi- 
nution of  vapor-pressure."  Raoult  formulated  this  rela- 
tion as  follows: 

"For  ethereal  solutions  the  relative  diminution  of  vapor- 
pressure  is  proportional  to  the  number  of  molecules  of  non- 
volatile substance  in  100  molecules  of  the  solution." 

For  dilute  solutions,  K,  as  we  have  seen,  is  very  nearly 
1.0,  when  the  above  expression  becomes, 

/_  f 

=  0.01 


flf 

Raoult  tested  this  deduction  by  dissolving  a  number  of 
substances  hi  ether,  with  the  following  results.1 


flf 

Oil  of  turpentine  0.0099 

Methyl  salicylate  0.0094 

Benzoic  acid  0.0097 

Aniline  0.0106 

1  Ann.  Chim.  Phys.  [6],  16,  400  (1888). 


110  THE  NATURE  OF  SOLUTION 

Effect  of  Nature  of  Solvent  —  Generalization  of  Raoult.  — 
Raoult  had  thus  shown  that  the  lowering  of  the  vapor-ten- 
sion of  ether  is  dependent  only  on  the  ratio  between  the 
molecules  of  the  solvent  and  of  the  dissolved  substance,  and 
is  practically  independent  of  the  nature  of  the  dissolved 
substance.  There  still  remained,  however,  one  more  funda- 
mental question.  Is  the  relative  lowering  of  the  vapor- 
tension  dependent  upon,  or  independent  of  the  nature  of 
the  solvent? 

This  question  was  answered  very  satisfactorily  by  Raoult, 
in  another  paper.1  He  measured  the  lowering  of  the  vapor- 
tension  of  twelve  solvents  by  different  substances  dissolved 
in  them,  and  from  the  results  calculated  the  following 
relations. 

Let  M  be  the  molecular  weight  of  the  solvent  and  K  the 

TT 

molecular  lowering  of  its  vapor-pressure,  the   ratio  TT  is 

M 

the  relative  lowering  of  vapor-pressure  produced  by  one 
molecule  of  solute  in  one  hundred  molecules  of  the  solvent, 
regardless  of  the  nature  of  both  solute  and  solvent.  His 

results  for  a  few  solvents  are  given  below. 

K 

Solvent  M  K  M 

Water  18.  0.185  0.0102 

Carbon  disulphide  76.0  0.80  0.0105 

Chloroform  119.5  1.30  0.0109 

Benzene  78.0  0.83  0.0106 

Ether  74.0  0.71  0.0096 

Acetone  58.0  0.59  0.0101 

Methyl  alcohol  32.0  0.33  0.0103 

Raoult  points  out  that  although  the  values  oiK  and  of 
M  vary  as  much  as  one  to  nine,  the  ratio  between  the  two  is 
very  nearly  constant,  the  constant  having  the  approximate 
value  of  0.0105. 

Raoult  then  formulates  what  has  come  to  be  known  as 
"the  Law  of  Raoult,"  which  is  a  generalization  that  is 
independent  of  the  chemical  nature  both  of  the  solvent  and 
of  the  dissolved  substance.2 

1  Compt.  Rend.,  104,  1430  (1887).  2  Compt.  Rend.,  104,  1433  (1887). 


DEPRESSION  OF  THE  VAPOR-TENSION  111 

"A  molecule  of  any  non-volatile  solid  other  than  salts 
in  solution  in  one  hundred  molecules  of  any  volatile  liquid, 
produces  a  lowering  of  the  vapor-tension  of  the  solvent 
which  is  nearly  a  constant  part  of  the  value  of  this  quantity 
—  the  constant  being  very  nearly  0.0105." 

This  generalization,  which  we  now  know  holds  only  ap- 
proximately, enables  us,  as  Raoult  pointed  out,  to  use  this 
property  of  solutions  to  determine  the  molecular  weight  of 
the  dissolved  substance.  Another  method  based  upon  the 
same  principle  as  this,  is,  however,  so  far  to  be  preferred  that 
this  application  of  the  lowering  of  vapor-tension  will  only 
be  referred  to. 

Raoult  followed  one  lead  after  another,  first  studying 
the  effect  of  the  dissolved  substance  and  arriving  at  a  partial 
generalization;  then  the  effect  of  the  solvent,  and  com- 
pleting the  generalization  which  shows  that  lowering  of 
vapor-tension,  like  osmotic  pressure,  is  a  property  which 
depends  upon  numbers,  and  numbers  only  —  the  ratio 
between  the  number  of  parts  of  the  solvent  and  of  the  dis- 
solved substance.  This  work  of  Raoult,  from  the  standpoint 
of  method,  as  well  as  of  bearing  on  modern  chemical  thought, 
is  to  be  ranked  among  the  classics  of  chemistry. 

Method  of  Frazer  and  Lovelace  for  Measuring  Vapor- 
Tension. —  This  method1  applied  to  the  vapor-pressures  of 
aqueous  solutions  is  a  differential,  static  method.  It  consists 
in  measuring  directly  the  difference  between  the  vapor- 
tension  of  the  pure  solvent  and  that  of  the  solution,  by 
means  of  a  very  sensitive  manometer  designed  by  Lord  Ray- 
leigh  for  the  study  of  the  behavior  of  gases  at  low  pressure. 

The  study  of  the  vapor-pressure  of  solutions  by  the 
static  method  necessitated  special  precautions  to  eliminate 
error  due  to:  (1)  temperature  changes,  (2)  changes  in  the 
surface  concentration  of  the  solution,  and  (3)  incomplete 
removal  of  dissolved  gases  from  the  solvent  and  from  the 
solution.  The  first  source  of  error  is  practically  removed 

1  Journ.  Amer.  Chem.  Soc.,  36,  2439  (1914);  Zeit.  phys.  Chem.,  89,  155 
(1914). 


112 


THE  NATURE  OF  SOLUTION 


by  using  a  thermostat  capable  of  such  accurate  control, 
that  the  variations  in  temperature  cannot  be  detected  by  a 
Beckmann  thermometer  reading  to  0.001°  C. 

The  error  due  to  change  in  surface  concentration  is 
eliminated  by  vigorous  and  continuous  stirring  in  a  vacuum 
of  both  solution  and  solvent. 

Traces  of  dissolved  gases  are  removed  by  allowing  the 
vapor  from  the  solution  and  solvent  to  expand  into  an  ex- 


FIQ.  2. 

hausted  flask.  The  operation  is  repeated  again  and  again, 
allowing  sufficient  time  at  each  exhaustion  for  the  space  to 
become  saturated  with  the  dissolved  gases. 

To  avoid  the  leakage  of  air  into  the  apparatus  all  con- 
nections are  sealed  and  no  stopcocks  are  used. 

The  measurement  of  the  difference  in  the  vapor-pressure 
of  solution  and  of  solvent  is  made  by  means  of  the  Ray- 
leigh  manometer  shown  in  figure  2.  It  consists  of  a  Y-- 
shaped glass  tube.  On  each  of  the  upper  limbs  is  blown  a 
bulb  about  40  mm.  in  diameter.  The  lower  limb  of  the  Y 
tube  is  of  barometric  height  and  to  its  end  is  attached  a 


DEPRESSION  OF  THE  VAPOR-TENSION  113 

mercury  reservoir,  the  height  of  which  can  be  accurately 
adjusted  by  turning  the  screw  /  shown  to  the  right  of  the 
apparatus.  The  glass  Y  tube  is  set  with  plaster  of  Paris 
in  an  iron  pot  M .  M  is  pivoted  at  Z),  and  the  instrument 
may  be  rotated  on  D  by  means  of  the  screw  0. 

When  the  apparatus  is  exhausted  the  mercury  rises  into 
the  two  limbs  of  the  Y  tube.  By  proper  manipulation  of  / 
and  0,  the  two  ground-glass  points  RR  are  made  just  to 
touch  their  images  in  the  mercury,  coincidence  being  deter- 
mined by  means  of  two  microscopes  (not  shown  hi  the 
drawing).  At  CC  two  glass  tubes  PP  are  sealed.  One 
of  these  connects  one  limb  of  the  Y  with  a  glass  bulb  contain- 
ing air-free  solvent,  and  the  other  in  like  manner  connects 
the  other  limb  with  the  vessel  containing  the  solution.  The 
bulbs  containing  the  solution  and  solvent  are  immersed 
hi  the  thermostat  bath  mentioned  above. 

When  the  apparatus  is  completely  exhausted,  or  when  the 
same  pressure  exists  hi  the  two  limbs  of  the  Y  tube,  and 
the  adjustment  of  the  points  made  as  mentioned  above,  the 
instrument  is  in  the  zero  position,  the  points  being  exactly 
at  the  same  height.  When  there  is  a  difference  in  pressure 
on  the  two  sides,  as  when  one  of  the  tubes  P  communicates 
with  the  solution  and  the  other  with  the  pure  solvent,  the 
mercury  surface  is  depressed  more  in  the  limb  hi  which  the 
greater  pressure  is  exerted,  and  it  is  necessary  to  rotate 
the  instrument  by  means  of  0  a  certain  amount  in  order 
to  bring  both  points  RR  into  coincidence  again  with  then* 
images.  The  amount  of  this  rotation  from  the  zero  position 
is  determined  by  means  of  the  metallic  mirror  A,  mounted 
directly  on  the  instrument,  and  a  telescope  and  scale  mounted 
about  three  meters  in  front  of  A.  Having  determined  the 
amount  of  rotation  from  the  zero  position,  and  knowing 
accurately  the  distance  between  the  points  RR,  and  the 
distance  from  A  to  the  scale,  the  difference  in  the  level  of 
the  points  can  be  readily  calculated.  This  is  the  depression  of 
the  vapor-tension  of  the  solvent  by  the  dissolved  substance 
in  millimeters  of  mercury,  which  is  the  quantity  desired. 


114  THE  NATURE  OF  SOLUTION 

A  few  of  the  results  obtained  by  Frazer  and  Lovelace 
are  given  below. 

VAPOR-PRESSURE  LOWERINGS  OP  SOLUTIONS  OP  MANNITE  AT  20* 

Concentration  Observed  lowering 
0.1  normal  0.029  mm  Hg. 

0.2      "  0.059    "      " 

0.3      "  0.090    "      " 

0.4      "  0.118    "      " 

0.5       "  0.153     "      ^;- 

VAPOR-PRESSURE  LOWERINGS  OP  SOLUTIONS  OP  POTASSIUM  CHLORIDE  AT  20° 

Concentration  Observed  lowering 

0.2  normal  0.110  mm.  Hg. 

0.4      "  0.217    "      " 

0.6      "  0.329     "      " 

1.0      "  0.547    "      " 

1.5       "  0.826    "      " 

2.0      "  1.102     "      " 

Other  investigators  have  recently  developed  methods 
for  measuring  the  vapor-tensions  of  solvents  and  of  solu- 
tions, but  reference1  only  can  be  made  to  their  work. 

The  Boiling-Point  Method.  —  The  vapor-tension  of 
most  solvents  at  ordinary  temperatures  is  small  and  the 
lowering  of  this  tension  by  dissolved  substances  is  a  small 
part  of  this  small  quantity,  as  we  have  seen.  The  measure- 
ment of  a  very  small  quantity  is  always  difficult,  and  unless 
some  very  refined  method  such  as  that  of  Frazer  and  Love- 
lace is  employed,  the  experimental  error  is  relatively  large. 
For  general  practical  use  in  the  laboratory  such  methods  are 
not  adapted;  and  it  is  important  to  know  the  molecular 
weights  of  substances  in  general,  hi  solvents  in  general. 
For  this  purpose  the  vapor-tension  method  has  been  aban- 
doned in  favor  of  a  method  based  upon  the  same  principle, 
but  which  is  incomparably  simpler  to  apply. 

The  vapor-tension  method  depends  upon  heating  the 
solution  and  the  solvent  to  the  same  temperature,  and 
measuring  the  vapor-tensions  of  the  two  at  the  constant 
temperature. 

i  Tower:  Journ  Amer.  Chem.  Soc.,  30, 1219  (1908);  Miindel:  Zeit.  phys. 
Chem.,  85,  435  (1913);  Washburn  and  Heuse:  Journ.  Amer.  Chem.  Soc., 
37,  309  (1915^. 


DEPRESSION  OF  THE  VAPOR-TENSION  115 

Instead  of  keeping  the  temperature  constant  and  meas- 
uring the  different  vapor-pressures  at  this  constant  tem- 
perature, we  can  heat  the  solution  and  the  solvent  until 
the  vapor-pressures  of  the  two  are  equal,  and  then  measure 
the  temperatures  required  to  bring  about  this  condition. 
The  constant  pressure  chosen  is  the  pressure  of  the  atmos- 
phere, and  the  temperatures  are  the  boiling-points  of  so- 
lution and  solvent;  i.e.,  the  temperatures  necessary  to 
produce  vapor  over  both  solution  and  solvent  which  will 
just  overcome  the  atmospheric  pressures.  The  vapor-tension 
method  thus  becomes  the  "  boiling-point  method." 

Boiling-point  Method  of  Beckmann.  —  The  first  to  de- 
vise a  reasonably  satisfactory  method  for  determining  the 
boiling-points  of  solutions  and  solvents  was  Beckmann. 
The  essential  features  of  any  good  boiling-point  method  are 
a  very  sensitive  thermometer,  and  good  heat  insulation 
around  the  boiling  liquid.  The  thermometer  devised  by 
Beckmann  has  a  very  large  bulb,  and  also  a  very  fine 
capillary,  but  the  peculiar  feature  of  it  is  at  the  top.  It  is 
provided  here  with  a  reservoir.  By  proper  manipulation, 
more  or  less  of  the  mercury  from  the  bulb  may  be  trans- 
ferred to  this  reservoir  so  as  to  adjust  the  reading  of  the 
thermometer  on  the  scale  for  widely  different  temperatures. 
The  Beckmann  thermometer  is  therefore  a  purely  (Inferential 
thermometer. 

Beckmann1  has  designed  a  large  number  of  forms  of 
apparatus  for  determining  the  boiling-points  of  solutions 
and  of  solvents.  All  are  based  upon  the  same  principle. 
The  liquid  whose  boiling-point  is  to  be  determined  is  placed 
in  a  glass  vessel  of  test-tube  shape,  and  this  is  surrounded  by 
an  asbestos  jacket  or  a  vessel  containing  some  of  the  same 
liquid  heated  to  boiling.  The  thermometer  is  plunged 
directly  into  the  liquid.  All  of  these  forms  appear  to  the 
writer  to  have  one  serious  defect  —  they  do  not  adequately 
prevent  the  cold,  recondensed  solvent  from  coming  hi  con- 

1  See  the  Author's  Elements  of  Physical  Chemistry,  4th  edition  (The 
Macmillan  Co.),  p.  263. 


116 


THE  NATURE  OF  SOLUTION 


8 


tact  with  the  thermometer  before  it  has  been  reheated  to  its 

true  boiling-point,  and  they 
do  not  sufficiently  safeguard 
the  apparatus  and  thermometer 
from  the  effects  of  radiation. 
To  overcome,  as  far  as  possible, 
these  objections  the  following 
form  of  boiling-point  apparatus 
was  designed  by  the  author. 

Boiling-Point  Apparatus  of 
Jones.  —  The  apparatus  de- 
signed by  Jones1  is  sketched  in 
figure  3.  It  is  simply  a  glass 
tube  with  glass  beads  and  plati- 
num scraps  in  the  bottom  and 
a  condenser  attached;  a  plati- 
num cylinder  D  being  intro- 
duced as  shown  in  the  figure; 
the  object  of  the  metal  cylin- 
der being,  in  the  first  place,  to 
cut  off  the  colder,  recondensed 
solvent  from  the  bulb  of  the 
thermometer,  until  the  liquid 
has  been  heated  again  to  the 
boiling-point.  This  colder  liq- 
uid must  pass  down  through 
the  entire  length  of  the  boiling 
liquid  before  it  can  enter  the 
cylinder  from  below,  and,  con- 
sequently, is  heated  again  to 
the  boiling-point. 

The  platinum  cylinder  serves 
also  to  reduce  to  a  minimum  the 
effect  of  direct  radiation.   If  the 
warm  bulb  of  the  thermometer 
FIG.  3.  is  not  surrounded  completely  by 

i  Amer.  Chem.  Journ.,  19,  581  (1897). 


DEPRESSION  OF  THE  VAPOR-TENSION  117 

metal,  it  will  radiate  heat  outward  through  the  liquid  to 
colder  objects  in  the  neighborhood;  liquids  being  some- 
what transparent  to  the  longer  heat  rays.  This  effect  is  re- 
duced to  a  minimum  by  surrounding  the  thermometer  bulb 
with  metal  heated  to  the  same  temperature  as  the  bulb  of 
the  thermometer  itself.  On  top  of  the  beads  in  the  bottom 
of  the  glass  vessel  a  few  scraps  of  platinum  are  placed,  to 
prevent  the  radiation  of  heat  directly  from  the  flame  to  the 
bulb  of  the  thermometer. 

Beckmann,  in  his  later  forms  of  apparatus,  surrounded 
the  tube  containing  the  liquid  to  be  boiled  with  a  double- 
walled  vessel  containing  between  its  two  walls  some  of  the 
same  liquid  whose  boiling-point  was  to  be  determined. 
This  was  boiled  at  the  same  tune  that  the  liquid  hi  the  inner 
cylinder  was  boiled.  In  this  way  the  effect  of  radiation 
was  somewhat  diminished,  but  by  no  means  rendered 
negligible.  The  platinum  cylinder  accomplishes  both  of  the 
purposes  for  which  it  was  introduced,  as  is  shown  by  the 
results  which  have  been  obtained  with  this  form  of  ap- 
paratus. 

Molecular  Weights  from  Boiling-Point  Determinations. 
-  It  is  obvious  that  any  property  which  depends  only  on 
the  ratio  between  the  number  of  parts  of  solute  and  solvent 
—  an  arithmetical  property  —  is  independent  of  the  chemical 
nature  of  either  and  can  be  used  to  determine  the  molecular 
weights  of  substances  in  general,  dissolved  hi  solvents  hi 
general. 

A  large  amount  of  work  has  been  done  on  the  molecu- 
lar weights  of  a  great  variety  of  substances  hi  a  large  number 
of  solvents.  The  results  of  this  work  can  be  easily  sum- 
marized. Substances  in  general  hi  solvents  in  general  are 
hi  the  simplest  molecular  condition  —  i.e.,  dissolved  sub- 
stances are  in  the  same  state  of  aggregation  as  gases.  This 
doubtless  has  something  to  do  with  the  relations  discussed 
earlier  between  solutions  and  gases. 

There  are,  however,  many  exceptions  known  to  the  above 
general  relation.  Certain  substances  in  certain  liquids 


118  THE  NATURE  OF  SOLUTION 

are  polymerized,  and  in  some  cases  very  much  polymerized, 
just  as  certain  elements  and  compounds  in  the  form  of  va- 
por are  polymerized.  Acetone  is  a  solvent  which  has  very 
marked  polymerizing  power,  and  there  are  many  other  ex- 
amples of  a  similar  nature.  It  should  be  noted  that  there 
is  no  relation  known  between  the  molecular  weight  of  a  dis- 
solved substance,  and  that  of  the  same  substance  in  the 
pure,  homogeneous  state. 

Electrolytic  Dissociation  Measured  by  the  Boiling- 
Point  Method.  —  Another  application  of  the  boiling-point 
method,  which  in  some  respects  is  more  important  than  the 
problem  of  molecular  weights  in  solution,  is  the  measure- 
ment of  electrolytic  dissociation  in  nonaqueous  solvents. 
This  was  at  one  tune  very  important,  because  we  had  then 
absolutely  no  method  for  measuring  electrolytic  dissociation 
in  a  large  number  of  solvents,  other  than  the  boiling-point 
method,  for  reasons  which  will  appear  in  the  proper  place. 
The  boiling-point  method  of  Beckmann  could  not  be  used  for 
this  purpose,  because  it  was  too  inaccurate.  An  improved 
boiling-point  method  has  been  applied  to  this  problem. 

Jones,1  using  the  apparatus  designed  by  himself,  has 
measured  the  dissociation  of  several  salts  in  methyl  and  in 
ethyl  alcohols.  Theoretically  the  matter  is  very  simple. 
From  the  rise  in  the  boiling-point  produced  by  dissolving  a 
known  weight  of  substance  in  a  known  weight  of  solvent,  the 
rise  produced  by  dissolving  a  gram-molecular  weight  of  the 
substance  in  1000  grams  of  the  solvent  was  calculated.  This 
is  known  as  the  "molecular  rise."  The  molecular  rise  for 
the  substance  in  question  divided  by  the  constant  for  the 
solvent  in  question  gives  what  is  known  as  the  Van't  Hoff 
coefficient  "  i"  since  it  is  the  coefficient  which  was  intro- 
duced into  the  simple  gas  equation  PV  =  RT  to  make  it  apply 
to  the  osmotic  pressures  of  electrolytes. 

The  dissociation  X  for  binary  electrolytes  or  those  which 
break  down  into  two  ions  each  is,2  X  =  i  —  1. 

1  Zeit.  phys.  Chem.,  31,  114  (1899). 

2  Ibid.,  11,  110,  529;  12,  639  (1893). 


DEPRESSION  OF  THE  VAPOR-TENSION  119 

For  ternary  electrolytes  or  those  whose  molecules  break 

i  —  1 
down  into  three  ions  each,  X  =  — r—  and  so  on. 

A  few  of  the  results  obtained  hi  this  laboratory1  are 
given  hi  the  following  table. 

Dissociation  Dissociation 

Cone  in  methyl  alcohol  in  ethyl  alcohol 

KI                              0.1  52.0%  25.0% 

Nal                            0.1  60.0  "  33.0  " 

NHJ                         0.1  50.0  " 

KBr                           0.1  50.0  " 

NH4Br                       0.2  49-0  "  21.0  " 

CHsCOONa              0.1  38.0  "  14.0  " 

These  results  are,  of  course,  to  be  regarded  only  as 
approximations,  the  error  in  the  most  refined  boiling-point 
method  being  considerable,  if  for  no  other  reason,  because 
the  boiling-point  of  a  solvent  is  affected  so  markedly  by 
slight  changes  in  the  barometer. 

Molecular  Weights  of  the  Metals  in  Mercury.  —  Before 
leaving  the  subject  of  vapor- tension,  there  is  one  other  ap- 
plication of  this  method  of  detennining  molecular  weights 
which  must  be  considered. 

Ramsay2  in  1899  studied  the  lowering  of  the  vapor-ten- 
sion of  mercury  by  metals  dissolved  in  it.  Knowing  the 
vapor-tension  constant  of  mercury  and  the  concentration 
of  the  amalgam  in  question,  he  could  calculate  the  molecular 
weight  of  the  metal  hi  question  in  the  mercury,  under  the 
conditions  of  the  experiment.  A  few  of  his  results,  taken 
from  the  paper  referred  to  above,  are  given. 

Atomic 
weight 
7.02 
23.04 
23.04 
23.04 
39.14 
39.14 
40.08 
137.00 
24.30 
24.30 
55.00 
197.22 
197.22 

Taken  from  Ibid.,  31,  140  (1899).       *  Journ.  Chem.  Soc.,  65,  521  (1889). 


Number  of  atoms 

Molecular 

Metal 

per  100  atoms  Hg. 

weight 

Li 

1.70 

7.1 

Na 

0.86 

21.6 

Na 

1.87 

18.3 

Na 

5.35 

15.1 

K 

1.55 

29.1 

K 

5.26 

30.2 

Ca 

0.19 

19.1 

Ba 

0.90 

75.7 

Mg 

0.70 

24.0 

Mg 

4.82 

21.5 

Mn 

1.14 

55.5 

Au 

1.59 

207.4 

Au 

2.80 

208.1 

120  THE  NATURE  OF  SOLUTION 

The  above  results  contain  a  number  of  points  of  interest. 
Most  of  the  metals  dissolved  in  mercury  are  in  the  simplest 
atomic  condition.  There  are,  however,  some  exceptions, 
and  these  are  the  interesting  features. 

The  molecular  weight  of  sodium  is  less  than  its  atomic 
weight,  and  this  becomes  still  smaller  as  the  concentration 
of  the  sodium  is  increased. 

The  most  remarkable  results  in  mercury  as  the  solvent 
were  obtained  with  calcium  and  barium.  When  dissolved 
in  mercury  their  molecular  weights  are  almost  exactly  half 
of  their  atomic  weights.  If  we  consider  that  the  amalgams  of 
these  metals  studied  by  Ramsay  were  very  dilute,  and  the 
experimental  error  therefore  larger,  we  would  not  be  inclined 
to  lay  so  much  stress  upon  these  data,  had  they  not  to  a  large 
extent  been  confirmed  by  subsequent  results  by  an  entirely 
different  method. 

Humphrey  and  Mohler,1  working  with  Rowland,  studied 
the  displacement  of  the  spectrum  lines,  when  the  incan- 
descent, elementary  gases  producing  them  were  under  pres- 
sure. They  found  that  from  the2  "product  of  the  cube  root 
of  the  ' atomic  volume'  and  the  coefficient  of  linear  expansion 
of  the  substance  in  the  solid  form,  certain  numbers  were  ob- 
tained whose  ratios  were  the  same  as  those  of  the  shift  for 
the  respective  elements." 

The  atomic  volume  is  the  atomic  weight  divided  by 
the  specific  gravity  or  density  in  the  solid  state.  When 
the  displacement  of  the  calcium  lines  were  measured,  and  the 
results  compared  with  the  displacement  calculated  from  the 
above  relation,  the  two  would  agree  for  calcium  only  when  the 
assumption  was  made  that  the  atomic  weight  of  this  element 
in  its  highly  heated  vapor  was  not  40,  but  a  smaller  value. 

The  same  result  was  then  reached  by  Ramsay  when 
certain  metals  are  dissolved  in  mercury  at  comparatively 
low  temperatures,  and  by  Humphreys  and  Mohler,  when  the 
metals  are  volatilized  in  the  electric  arc. 

1  Astrophysical  Journal,  3,  114  (1896). 

2  Ibid.,  3,  131  (1896). 


DEPRESSION  OF  THE  VAPOR-TENSION  121 

At  the  time  that  these  observations  were  made  they  were 
very  surprising.  We  had  been  accustomed  to  think  of  the 
chemical  atoms  as  ultimate  units  which  could  not  be  broken 
down  into  anything  simpler.  Now  hi  the  light  of  the  work 
of  Thomson  this  conclusion  is  not  so  surprising.  We  know 
that  the  atoms  are  complex,  and  there  is  no  a  priori  reason 
why  they  should  not  yield  simpler  things.  Indeed,  electrons 
have  been  obtained  from  many  of  them. 


CHAPTER  VIII 

DEPRESSION   OF   THE   FREEZING-POINT   OF   A    SOLVENT 
BY  THE  SOLUTE 

THAT  dissolved  substances  lower  the  freezing-point  of  the 
solvent  has  been  known  qualitatively  for  a  very  long  tune. 
The  first  to  have  made  quantitative  measurements  in  this  field 
seems  to  have  been  Blagden.1  As  early  as  1788,  he  found  that 
the  amount  of  the  lowering  of  the  freezing-point  is  propor- 
tional to  the  amount  of  dissolved  substance  in  a  given 
volume  of  the  solution. 

Investigations  of  Raoult.  —  Very  little  real  progress  was 
made  in  the  study  of  freezing-point  lowering  from  the  tune 
of  Blagden,  until  the  problem  was  taken  up  by  the  one 
who,  we  have  seen,  did  such  magnificent  work  in  connec- 
tion with  the  lowering  of  vapor-tension  of  solvent  by  dis- 
solved substance  —  Raoult. 

Coppet,2  nearly  one  hundred  years  after  Blagden,  did 
make  one  advance  of  some  consequence  hi  the  study  of 
freezing-point  lowering,  in  that  he  expressed  his  concentra- 
tions in  terms  of  molecular  quantities,  and  thus  made  the 
results  with  different  substances  comparable  with  one 
another.  Coppet  consequently  discovered  that  correlated 
substances  produce  essentially  the  same  lowering  of  the 
freezing-point  of  water. 

Raoult3  attacked  the  problem  of  lowering  of  freezing-point 
in  much  the  same  way  that  he  had  attacked  the  question  of 
lowering  of  vapor-tension.  Instead  of  limiting  his  investi- 
gations to  water  as  a  solvent,  he  extended  them  to  a  fairly 
l^rge  number  of  solvents — acetic  acid,  formic  acid,  ben- 

1  Phil.  Trans.,  78,  277  (1788). 

2  Ann.  Chim.  Phys.  [4],  23,  366  (1871);  25,  502  (1872);  26,  98  (1872). 
«  Ibid.  [5],  28,  133  (1883);  [6],  2,  66  (1884). 


DEPRESSION  OF  THE  FREEZING-POINT  123 

zene,  nitrobenzene,  ethylene  bromide.  This  was  obviously 
the  scientific  way  to  proceed.  Otherwise,  relations  holding 
for  one  or  for  several  solvents  might  be  discovered  which 
would  have  no  significance  when  still  other  solvents  were 
brought  within  the  scope  of  the  work.  A  few  of  the  results 
obtained  by  Raoult  hi  the  different  solvents  will  be  given. 

Raoult  called  the  lowering  of  the  freezing-point  pro- 
duced by  one  gram  of  substance  hi  one  hundred  grams  of 
solvent,  the  "coefficient  of  lowering "  of  the  substance  hi 
question.  This  he  represented  by  A.  M  is  the  molecular 
weight  of  the  dissolved  substance,  and  T  the  molecular 
lowering  of  the  freezing-point.  We  have, 

M  A  =  T 


RESULTS  OBTAINED  BY  RAOULT  l 
SOLVENT  ACETIC  Aero 

Molecular 

Substance  lowering 

dissolved                                    Formula  M A  -  T 

Chloroform                                          CHCU  38.6 

Carbon  tetrachloride                          CC14  38.9 

Naphthalene                                       CioHs  39.2 

Aldehyde                                             C,H4O  38.4 

Benzoic  acid                                        CrHeOj  43.0 

Ethyl  alcohol                                      CjHeO  36.4 

Hydrochloric  acid                               HC1  17.2 

Sulphuric  acid                                     HjSO4  18.6 

From  the  results  of  the  study  of  about  sixty  compounds 
hi  this  solvent,  Raoult  concluded  that  the  molecular  lower- 
ings  in  acetic  acid  approach  two  numbers  —  39  and  18;  the 
one  being  approximately  the  half  of  the  other. 

In  formic  acid  the  following  results  were  obtained: 

Molecular 

Substance  lowering 

dissolved  Formula  MA  -  T 

Chloroform  CHCla  26.5 

Benzene  C6H6  29.4 

Ether  C4HioO  28.2 

Acetic  acid  C2H4O2  26.5 

Arsenic  trichloride  AsCl3  26.6 

Magnesium  formate  (HCO*)2Mg  13.9 

1  Taken  from  Ann.  Chim.  Phys.  [6],  2,  71  to  84  (1884). 


124  THE  NATURE  OF  SOLUTION 

The  molecular  lowerings  in  formic  acid  approach  the  two 
values  28  and  14. 

The  molecular  lowerings  in  benzene  are  grouped  around 
49  and  25.  Similarly,  the  molecular  lowerings  in  nitro- 
benzene group  themselves  around  73  and  36;  and  in  ethyl- 
ene  bromide  around  118  and  58. 

A  few  of  the  results  in  nitrobenzene  are  given. 

Molecular 

Substance  lowering 

dissolved  Formula  M A  =  T 

Chloroform  CHC18  69.9 

Carbon  disulphide  CS2  70.2 

Naphthalene  CioHs  73.6 

Ether  C4HioO  67.4 

Acetone  C3H6O  69.2 

Stannic  chloride  SnCl4  ,     71.4 

Methyl  alcohol  CH4O  35.4 

Ethyl  alcohol  CzKUO  35.6 

Acetic  acid  C2H4O2  36.1 

Here  again  we  have  the  two  sets  of  values,  one  of  which 
is  just  about  half  the  other. 

Raoult  studied  about  one  hundred  compounds  in  water, 
and  a  few  of  his  results  for  aqueous  solutions  are  given. 

Molecular 

Substance  lowering 

dissolved  Formula  M  A  =  T 

Hydrobromic  acid  HBr  39.6 

Phosphoric  acid  H3PO4  42.9 

Sodium  hydroxide  NaOH  36.2 

Potassium  chloride  KC1  33.6 

Sodium  nitrate  NaNO3  34.0 

Calcium  hydroxide  Ca(OH)2  48.0 

Calcium  chloride  CaCl2  49.9 

Strontium  nitrate  Sr(NO3)2  41.2 

Sulphurous  acid  H2SO3  20.0 

Boric  acid  H3BO3  20.5 

Methyl  alcohol  CH4O  17.3 

Dextrose  C6Hi2O6  19.3 

Formic  acid  H2CO2  19.3 

Urea  CON2H4  17.2 

The  values  for  water  produced  by  compounds  which  are 
dissociated  are  around  37;  the  other  class  of  values  is 
around  18.5. 

What  is  the  meaning  of  the  two  values  in  any  given  sol- 
vent? Raoult  points  out  that  in  a  constant  weight  of  any 


DEPRESSION  OF  THE  FREEZING-POINT  125 

solvent  all  of  the  physical  molecules,  independent  of  their 
nature,  produce  the  same  lowering  of  the  freezing-point. 
He  attempts  to  interpret  the  two  values  found  for  every 
solvent  hi  the  light  of  this  idea.  When  the  maximum 
molecular  lowering  is  produced  it  means  that  the  physical 
molecule  contains  one  of  the  simplest  chemical  molecules 
of  the  substance  —  the  two  are  identical.  When  some  of 
the  chemical  molecules  are  united  hi  pairs,  the  molecular 
lowering  is  less  than  the  maximum.  When  all  the  chemical 
molecules  are  united  in  pairs,  the  molecular  lowering  is 
just  half  the  maximum.  The  molecular  lowering  of  water  is 
sometimes  kss  than  half  the  maximum.  This  means,  accord- 
ing to  Raoult,  that  the  chemical  molecules  are,  hi  part  at 
least,  aggregated  in  groups  of  three  to  form  the  physical 
molecules,  which  is  the  unit  in  determining  the  lowering  of 
the  freezing-point  of  a  solvent. 

Raoult  explains  thus  the  results  with  different  compounds 
in  the  same  solvent,  but  this  does  not  explain  the  results 
with  the  same  compound  in  different  solvents;  which  is,  of 
course,  necessary  before  any  comprehensive  generalization 
can  be  reached. 

Before  this  can  be  done,  we  must  calculate  the  results 
in  the  different  solvents  on  a  comparable  basis.  This  is 
most  easily  accomplished  by  referring  the  results  to  one 
molecule  of  solute  in  100  molecules  of  solvent.  To  do  this 
it  is  only  necessary  to  divide  the  molecular  lowering  for 
each  substance  T,  by  the  molecular  weight  of  the  sol- 
vent M. 

If  we  represent  the  lowering  produced  by  one  mole- 
cule of  substance  in  100  molecules  of  solvent,  by  T',  we  have 
• 

T 

T'=  — 
M 

In  the  following  table  are  the  results  for  six  solvents, 
T  being  the  maximum  molecular  lowering  of  the  freezing- 
point.1 

1  Taken  from  Ann.  Chim.  Phys.  [6],  2,  91  (1884). 


126  THE  NATURE  OF  SOLUTION 

Solvent  M  T  ~ 

Water  18  47  2.61° 

Formic  acid  46  29  0.63 

Acetic  acid  60  39  0.65 

Benzene  78  60  0.64 

Nitrobenzene  123  73  0.59 

Ethylenebromide  188  119  0.63 

Omitting  water,  the  ratio  T'  is  very  nearly  constant,  vary- 
ing only  between  0.59°  and  0.65°  — the  mean  being  0.63°. 

Raoult  would  explain  the  apparent  abnormality  presented 
by  water  as  follows.  If  we  assume  that  the  physical  mole- 
cule of  water  is  made  up  of  three  of  the  simplest  chemical 
molecules,  we  can  explain  all  of  the  molecular  lowerings  of 
this  solvent,  which  are  in  the  neighborhood  of  37. 

37 
18^3  =0'685' 

More  recent  work,  however,  gave  molecular  lowerings  of 
water  which  are  as  great  as  47.  To  account  for  this  value, 
we  must  assume  that  the  physical  molecule  of  water  is  made 
up  of  four  of  the  simplest  chemical  molecules. 

1^4=  °'65  '      '' 

which  is  very  nearly  the  mean  for  the  other  solvents  —  0.63. 

It  is  interesting  to  note  that  this  assumption  of  Raoult, 
that  the  molecular  weight  of  water  at  zero  degrees  is  (H20)4, 
has  subsequently  been  confirmed  by  the  work  of  Ramsay 
and  Shields,1  on  the  molecular  weights  of  pure  liquids.  They 
found  that  the  formula  of  water  at_and  near  zero  degrees 
is  H804  =  (H20)4. 

Raoult' s  Law  of  Freezing-Point  Lowering.  —  Explaining 
the  apparent  abnormalities  in  freezing-point  lowering,  as 
Raoult  did,  as  due  on  the  one  hand  to  polymerization  of  the 
molecules  of  the  solute,  and  on  the  other  to  polymerization 
of  molecules  of  the  solvent,  he  was  able  to  work  out  his  now 
famous  general  law  of  freezing-point  lowering. 

1  Zeit.  phys.  Chem.,  12,  433  (1893). 


DEPRESSION  OF  THE  FREEZING-POINT 


127 


One  molecule  of  any  substance  dissolved  in  one  hundred 
molecules  of  any  solvent  lowers  the  freezing-point  of  the  solvent 
nearly  a  constant  amount,  which  is  about  0.63°. 

This  law,  as  we  shall  see,  lies  at  the  basis  of  all  subse- 
quent work  done  hi  this  field.  This 
investigation  by  Raoult,  like  the  corre- 
sponding one  hi  the  field  of  the  lower- 
ing of  vapor-tension,  has  become  one  of 
the  classics  of  modern  chemistry. 

The  Freezing-Point  Method  of 
Beckmann.  —  The  first  to  devise  a 
fairly  satisfactory  method  for  measur- 
ing the  freezing-points  of  solvents  and 
ofc  solutions  was  Beckmann,  who,  it 
will  be  recalled,  also  devised  the  first 
reasonably  satisfactory  boiling-point 
method. 

The  key  to  the  freezing-point,  as  to 
the  boiling-point  method  of  Beckmann, 
is  his  thermometer,  which  has  already 
been  described  (p.  115).  The  remain- 
der of  his  apparatus  consists  of  two 
test-tubes,  one  placed  within  the  other 
—  the  air-space  between  the  two  being 
from  one  to  two  centimeters  wide.  The 
Beckmann  thermometer  is  inserted  into 
the  innermost  tube,  and  the  whole 
placed  in  a  suitable  freezing-mixture. 
The  liquid  around  the  thermometer  is 
provided  with  a  convenient  stirrer.  The  air-space  between 
the  two  test-tubes  is  for  the  purpose  of  protecting  the  liquid 
to  be  frozen  and  the  thermometer  immersed  hi  it,  from  a 
too  rapid  lowering  of  temperature  and  from  sudden  changes 
hi  temperature.  With  this  apparatus,  a  sketch  of  which  is 
given  in  figure  4,  Beckmann  determined  the  molecular 
weights  of  a  large  number  of  substances  in  a  fairly  large 
number  of  solvents. 


FIG.  4 


128  THE  NATURE  OF  SOLUTION 

The  freezing-point  constant  of  a  solvent  is  the  lower- 
ing of  its  freezing-point  in  degrees,  produced  by  dissolving 
hi  1000  grams  of  the  solvent  in  question  a  gram-molec- 
ular weight  *  of  a  completely  unpolymerized,  undissociated 
and  unhydrated  substance.  The  necessity  of  this  last- 
named  condition  will  appear  in  the  last  two  chapters  of  this 
book. 

The  freezing-point  constant  of  a  solvent,  like  its  boiling- 
point  constant,  can  be  calculated  from  the  equation, 

' 


100L 

which  was  deduced  by  Van't  Hoff2  and  tested  by  experi- 
ment. T  is  the  absolute  temperature  at  which  the  solvent 
freezes,  and  L  the  heat  of  fusion  of  one  gram  of  the  solvent. 

When  this  equation  is  used  to  calculate  the  boiling-point 
constant  of  a  solvent,  we  must  simply  reinterpret  the  symbols. 
T  is  then  the  absolute  temperature  at  which  the  solvent 
boils,  and  L  the  heat  of  vaporization  of  one  gram  of  the 
solvent. 

Results  of  Molecular  Weight  Determinations  by  the 
Freezing-Point  Method.  —  We  have  seen  that  the  molecular 
weights  of  substances  in  general  in  most  solvents,  at  the 
boiling-points  of  the  solvents,  are  the  simplest  possible. 
It  was  pointed  out  that  there  are  many  exceptions  to  this 
conclusion. 

The  same  general  conclusion  applies  to  the  molecular 
weights  of  substances  in  solvents  at  their  freezing-points  — 
they  are  hi  the  simplest  molecular  condition.  Here  again 
exceptions  present  themselves.  A  much  larger  number  of 
solvents  can  be  studied  by  the  boiling-point  than  by  the 
freezing-point  method.  More  solvents  boil  between  room 

1  Gram-molecular  weight  is  the  weight  in  grams  equal  to  the  molecular 
weight. 

2  Any  one  interested  in  the  deduction  of  this  equation  can  find  it  in  the 
Zeilschrift  fur  physikalische  Chemie,  1,  497  (1887),  or  somewhat  elaborated 
in  the  Author's  Elements  of  Physical  Chemistry,  4th  edition,  p.   254  (The 
Macmillan  Co.). 


DEPRESSION  OF  THE  FREEZING-POINT  129 

temperature  and,  say,  200°  than  freeze  near  ordinary  room 
temperature.  Nevertheless,  a  fairly  large  number  of  liquids 
do  freeze  at  temperatures  which  can  be  measured  with 
reasonable  accuracy.  Therefore,  the  above  conclusion  is 
based  upon  a  fair  amount  of  data. 

Substances  in  solution,  then,  both  at  the  freezing  and 
boiling-points  of  liquid  solvents,  are  hi  general,  hi  the  sim- 
plest molecular  condition,  i.e.,  in  the  same  condition  as  in 
the  gaseous  or  vapor  state.  As  has  been  pointed  out,  this 
probably  has  much  to  do  with  the  relations  between  solu- 
tions and  gases  discovered  by  Van't  Hoff,  and  discussed 
in  Chapter  V  of  this  book. 

Electrolytic  Dissociation  Measured  by  the  Freezing- 
Point  Method.  —  We  saw  when  studying  the  boiling-point 
method  that  its  most  important  application  was  to  the  prob- 
lem of  electrolytic  dissociation.  This  is  also  the  case  with 
the  freezing-point  method,  especially  when  historically  con- 
sidered. About  twenty-five  years  ago,  when  the  freezing- 
point  method  was  first  applied  to  the  problem  of  measuring 
electrolytic  dissociation,  there  were  only  two  other  methods 
known  for  dealing  with  this  problem.  One  of  these  was 
based  on  the  power  of  solutions  to  conduct  the  electric 
current  —  the  so-called  conductivity  method  —  and  with 
this  we  shall  become  familiar  a  little  later;  and  the  other 
was  based  upon  certain  relations  which  were  worked  out  in 
connection  with  saturated  solutions,1  and  which  it  would 
lead  us  too  far  to  discuss  here. 

The  point  in  the  present  connection  is,  that  these  two 
methods  did  not  give  concordant  results  and  this  was  used, 
and  with  entire  justice,  by  the  opponents  of  the  dissociation 
theory  as  an  argument  against  that  theory  —  there  was 
no  reliable  method  known  of  measuring  electrolytic  dis- 
sociation, even  if  it  existed. 

In  1892  Ostwald  started  the  author  at  the  task  of  so 
improving  the  freezing-point  method  of  Beckmann  that  it 
could  be  used  to  measure  electrolytic  dissociation,  to  see 

1  Zeit.  phys.  Chem.,  4,  372  (1889). 


130  THE  NATURE  OF  SOLUTION 

whether  the  results  obtained  by  this  method  would  agree 
with  those  obtained  by  either  of  the  other  two  methods,  and 
if  so,  with  which  one. 

Improved  Freezing-Point  Method.  —  The  first  step  was 
to  secure  a  far  more  sensitive  and  accurate  thermometer 
than  had  been  used  by  Beckmann.  Ostwald  had  Goetze 
construct  a  thermometer  which  was  ten  tunes  more  sensitive 
than  any  that  had  ever  been  made  before  that  tune.  The 
thermometer,  which  was  of  the  Beckmann  type,  contained 
in  the  bulb  about  200  grams  of  mercury,  and  the  capillary 
was  very  fine.  The  entire  scale  covered  a  range  of  only 
0.6°  and  was  divided  into  hundredths  and  thousandths  of 
a  degree.  With  a  telescope  it  was  a  very  simple  matter  to 
read  to  a  ten-thousandth  of  a  degree. 

The  next  part  of  the  problem  was  so  to  enlarge  the  vessel 
which  was  to  hold  the  solution  that  it  would  contain  a  large 
volume  of  the  solution.  The  vessel  selected  held  just  a 
liter  of  the  solution. 

A  very  efficient  stirrer  was  designed  which  would  keep 
the  liquid  well  stirred. 

This  apparatus  was  found  to  be  very  satisfactory  when 
applied  to  the  problem  in  hand.1 

Results  of  Measurements  of  Dissociation  by  the  Freezing- 
Point  Method.  —  The  dissociations  of  a  fairly  large  number 
of  salts,  over  as  wide  a  range  of  dilution  as  was  possible,  was 
worked  out  by  Jones  in  the  laboratory  of  Ostwald.  It  was 
soon  seen  that  there  is,  in  general,  a  close  agreement  be- 
tween the  values  found  by  the  freezing-point  method  and 
those  obtained  from  conductivity  measurements.  This 
can  be  readily  seen  from  the  following  table  of  data.  In 
this  table  the  data  from  freezing-point  lowerings  are  not 
taken  from  the  measurements  made  by  Jones  in  1892-1893; 
but  from  subsequent  work  done  by  Pearce2  hi  this  labora- 
tory in  1907. 

1  Zeit.  phys.  Chem.,  11, 110,  529  (1893);  12,  639  (1893). 

2  Amer.  Chem.  Journ.,  38,  683  (1907). 


DEPRESSION  OF  THE  FREEZING-POINT 


131 


Cone                 Dissociation  from  Dissociation  from 

Salt                 normal  freezing-point  lowering  conductivity 

( 0.10                            76.35  74.35 

CaCla                  <0.05                            80.96  80.62 

( 0.01                            90.61  89.67 

( 0.10  81.46  74.17 

SrCl2  <0.05  82.65  78.08 

( 0.01  91.87  89.37 

C  0.10  87.68  73.61 

MgCl2  Jo.05  90.97  79.78 

( 0.01  97.10  90.90 

( 0.025  90.10  81.13 

Ca(N08)2  <0.05  79.27  76.00 

( 0.25  72.83  61.28 

( 0.025  90.27  80.36 

Sr(N03)2  ]0.05  79.16  74.92 

( 0.10  73.32  68.59 

( 0.02  94.90  85.12 

Mg(N03)2  ^  0.05  84.24  78.80 

(0.10  81.95  74.78 

These  results  suffice  to  show  that  dissociation  as  meas- 
ured by  freezing-point  lowering,  is  usually  a  little  larger  than 
as  measured  by  the  power  of  the  solutions  to  conduct  the 
electric  current.  The  meaning  of  this  will  be  discussed  later. 

The  point,  however,  which  it  is  desired  to  bring  out  by 
means  of  the  above  results,  is  the  general  agreement  between 
the  two  sets  of  results  obtained  by  the  above  two  methods. 
An  examination  of  the  table  will  show  that  this  is  unmis- 
takable. 

Since  dissociation  measured  by  freezing-point  lowering 
agrees  in  general  with  the  results  obtained  by  means  of  the 
conductivity  method,  the  former,  from  what  was  stated 
earlier,  must  differ  from  the  values  obtained  by  the  solu- 
bility method.  We  had  then  three  methods  of  measuring 
dissociation,  —  conductivity,  freezing-point  lowering  and 
solubility.  The  first  two  gave  concordant  results,  and  the 
third  results  which  differed  from  those  determined  by  the 
other  two  methods.  This  condition  was  obviously  still  far 
from  satisfactory. 

Since,  however,  two  methods  gave  concordant  results 
it  seemed  probable  that  there  must  be  some  error  hi  the 


132  THE  NATURE  OF  SOLUTION 

results  obtained  by  the  third  method.  This  method  was 
re-examined  from  the  theoretical  side,  and  some  of  the  solu- 
bility experiments  repeated,  with  the  result  that  it  was  found 
that  in  applying  the  solubility  method  to  the  measurement 
of  dissociation  an  assumption  had  been  made  which  was 
erroneous;  and  when  this  error  was  corrected,  the  solu- 
bility method  gave  dissociation  values  which  agreed  very 
satisfactorily  with  those  obtained  by  both  the  conductivity 
and  the  freezing-point  methods. 

This  point  which  had  given  so  much  trouble  was  thus 
satisfactorily  cleared  up;  all  three  methods  of  measuring 
dissociation  giving  concordant  results. 

The  Three  Fundamental  Properties  of  Solutions. — 
We  have  seen  that  the  lowering  of  the  vapor-tension  of  sol- 
vents by  dissolved  substances  obeys  the  law  of  Raoult. 
This  means  that  lowering  of  vapor-tension  is  one  of  those 
properties  which  depends  upon  numbers  and  numbers  only, 
independent  of  the  nature  of  the  dissolved  substance  and 
independent  of  the  nature  of  the  solvent. 

We  have  also  seen  that  osmotic  pressure  is  another  of 
these  numerical  properties.  Further,  Arrhenius  deduced 
the  relation  between  the  lowering  of  vapor-tension  and  os- 
motic pressure. 

The  lowering  of  freezing-point,  obeying  the  law  of 
Raoult,  is  also  a  property  which  depends  for  its  magnitude 
only  on  the  relation  between  the  number  of  molecules  of  the 
solvent  and  of  the  dissolved  substance —  it  is  an  arithmetical 
property. 

The  relation  between  freezing-point  lowering  and  osmotic 
pressure  of  solutions  has  been  worked  out  by  Van't  Hoff.1 
Indeed,  he  has  deduced  the  equation  for  calculating  the 
freezing-  and  boiling-point  constants  of  a  solvent,  by  com- 
bining equations  for  its  osmotic  pressure  and  for  the  lowering 
of  the  freezing-point. 

Thus,  all  three  of  these  fundamental  properties  of  solu- 
tions have  been  connected  mathematically;  and  these 

1  Zett.  phys.  Chem.,  1,  496,  497  (1887). 


DEPRESSION  OF  THE  FREEZING-POINT  133 

are  the  three  fundamental  properties  of  all  true  solutions. 
Given  a  system  which  to  all  external  appearances  is  a  true 
solution,  we  shall  see  that  it  may  or  may  not  be  a  true 
solution.  How  are  we  to  determine? 

It  is  only  necessary  to  see  whether  it  has  an  osmotic 
pressure  which  obeys  the  laws  of  gas-pressure;  whether 
its  lowering  of  vapor-tension  and  its  freezing-point  lowering 
conform  to  the  laws  of  Raoult.  If  such  is  the  case,  it  is 
a  true  solution.  If  not,  it  is  a  colloidal  solution,  a  colloidal 
suspension  or  a  mechanical  suspension,  all  of  which  will  be 
discussed  later. 

While  dealing  with  these  numerical  properties  it  should 
be  pointed  out  that  they  are  the  most  important  properties, 
just  because  they  do  depend  on  numbers  and  numbers  only. 
It  is  the  study  of  such  properties  which  yields  results  of 
general  value. 

Whenever  we  find  a  property  which  changes  with  every 
change  in  the  composition  of  the  substance,  and  changes 
with  every  change  in  its  constitution,  the  study  of  such  a 
property  can,  at  best,  lead  only  to  empirical  relations;  and 
empirical  relations  usually  disappear  as  soon  as  enough  facts 
bearing  upon  them  are  brought  to  light. 

These  numerical  properties  may  be  called  the  fundamen- 
tal properties,  and  it  is  the  investigation  of  these  constant 
properties  which  really  advances  the  science. 

Freezing  of  Saturated  Solutions.  —  Before  leaving  the 
subject  of  freezing-point  lowering,  reference  must  be  made 
to  a  condition  which  exists  in  certain  solutions  at  their 
freezing-point.  When  an  ordinary  solution  is  frozen  the 
pure  .solvent  separates  in  the  solid  phase.  If  the  solution 
is  just  saturated  at  its  freezing-point,  what  will  separate 
when  the  solution  freezes?  Obviously,  the  solid  which  sep- 
arates will  be  a  mixture  of  the  solid  phase  of  the  solvent 
and  of  the  solute,  having  the  same  composition  as  the  solution 
itself.  This  mixture,  having  a  definite  composition,  will 
continue  to  separate  as  long  as  the  solution  saturated  at  its 
freezing-point  is  made  to  freeze  by  keeping  it  in  the  freezing- 


134  THE  NATURE  OF  SOLUTION 

mixture.  Since  the  solid  which  separates  has  a  constant 
composition,  the  question  which  would  naturally  suggest 
itself  is,  is  this  simply  a  mixture,  or  is  it  a  chemical  com- 
pound? Constant  composition  is  a  criterion  of  chemical 
union,  and  it  was  earlier  supposed  that  it  was  not  only  a  nec- 
essary, but  a  sufficient,  condition  hi  determining  whether  any 
given  system  is,  or  is  not,  a  chemical  compound. 

These  systems,  formed  by  freezing  solutions  which  are  just 
saturated  at  their  freezing-points,  Guthrie 1  called  cryohydrates, 
"cryo-"  from  the  fact  that  they  are  formed  at  the  freezing 
temperatures,  and  "hydrates"  because  in  the  case  of  aqueous 
solutions  they  contain  water,  even  if  in  the  solid  state. 

From  their  constant  composition  Guthrie  supposed  these 
substances  to  be  definite  chemical  compounds. 

Offer2  proved  this  not  to  be  the  case,  and  hi  the  following 
manner.  If  these  cryohydrates  are  chemical  compounds, 
then,  when  they  are  formed,  there  must  be  a  thermal  change, 
heat  being  either  evolved  or  absorbed,  as  in  all  chemical  reac- 
tions. Offer  found  that  the  amount  of  heat  required  to  melt 
a  cryohydrate,  was  just  equal  to  the  amount  of  heat  required 
to  melt  the  ice  in  the  cryhoydrate  and  dissolve  the  salt  con- 
tained in  it;  thus  showing  that  there  was  no  heat  either 
evolved  or  absorbed  when  the  cryohydrate  was  formed,  and 
therefore  cryohydrates  are  not  chemical  compounds. 

This  same  conclusion  was  confirmed  by  the  following 
fact.  Ice  dissolves  in  alcohol  when  the  whole  system  is  kept 
below  zero.  Take  a  cryohydrate  which  contains  a  salt 
insoluble  in  alcohol,  and  treat  it  with  alcohol.  The  ice  will 
dissolve  and  leave  the  solid  salt  behind,  showing  that  they 
were  not  in  a  state  of  chemical  combination,  but  were  simply 
a  mechanical  mixture  containing  the  two  constituents  hi  a 
certain  definite  proportion. 

It  is  very  fortunate  for  chemistry  that  cryohydrates  are 
not  chemical  compounds.  If  they  were,  the  total  number 
of  chemical  compounds  would  be  greatly  increased.  Every 


Phil  Mag.  [4],  49, 1  (1875);  [5],  1,  49  and  2,  211  (1876). 
Ber.  Wien.  Akad.,  81,  II,  1058  (1880). 


DEPRESSION  OF  THE  FREEZING-POINT  135 

solution  of  every  substance  in  every  solvent,  just  saturated 
at  the  freezing-point  of  the  solution,  when  frozen,  would 
yield  a  cryohydrate,  and  chemistry  would  thus  be  burdened 
with  a  still  larger  number  of  facts. 

The  reason  that  chemistry  as  a  science  has  developed 
more  slowly  than  physics  is  due  hi  part  to  the  nature  of 
chemical  phenomena.  They  obey  the  law  of  multiple  pro- 
portions, which  is  to  say,  that  chemical  combination  takes 
place  in  steps  —  1  of  A  to  1  of  B;  1  of  A  to  2  of  B;  2  of  A 
to  1  of  B;  and  so  on.  Such  phenomena  cannot  be  expressed 
hi  continuous  curves,  and  we  cannot  deal  mathematically  with 
discontinuous,  as  we  can  with  continuous  phenomena.  This 
condition  has  held  back  the  application  of  mathematics  to 
chemistry,  and  has  kept  it  longer  in  the  condition  of  an  em- 
pirical branch  of  science.  The  transformation  of  chemistry 
from  empiricism  to  an  exact  branch  of  science  was  rendered 
more  difficult  also  by  the  large  number  of  chemical  facts  that 
were  on  record.  The  large  number  of  compounds  and  the 
large  number  of  facts,  at  that  tune  chiefly  disconnected  and 
meaningless,  which  were  known  about  them,  made  it  ex- 
tremely difficult  to  discover  the  fundamental  laws  underlying 
these  phenomena,  and  to  which  they  conform. 

Were  the  number  of  such  facts  to  be  indefinitely  increased 
by  showing  that  cryohydrates  are  definite  chemical  com- 
pounds, it  would  be  an  unfortunate  day  for  chemistry. 

Some  of  the  fundamental  generalizations  which  have 
transformed  chemistry,  in  part  at  least,  from  empiricism 
through  system  into  science,  have  already  been  considered. 
The  most  important  of  these  are  the  relations  between  solu- 
tions and  gases  pointed  out  by  Van't  Hoff,  and  the  theory 
of  electrolytic  dissociation  of  Arrhenius.  To  these  should 
be  added  the  law  of  mass  action  discovered  and  mathemati- 
cally formulated  by  Guldberg  and  Waage,  and  Faraday's 
law  as  the  basis  of  valence,  which  will  be  discussed  later. 

It  should  be  noted  that  these  advances  have  all  been 
made  by  the  application  of  physical  and  mathematical 
methods  to  the  problems  of  chemistry. 


CHAPTER  IX 

AQUEOUS    SOLUTIONS    OF   ACIDS,   BASES,   AND 
SALTS  —  ELECTROLYTES 

THREE  of  the  most  important  classes  of  chemical 
compounds  are  acids,  bases,  and  salts.  These  three  classes 
taken  together  constitute  the  large  group  known  as  eke- 
trolytes.  All  of  these  compounds  when  brought  into  the 
presence  of  water  or  other  dissociating  solvents  are  broken 
down  to  a  greater  or  less  extent  into  ions,  are  electrolytically 
dissociated.  What  are  the  ions  of  acids  and  bases? 

What  is  an  Acid?  —  An  acid  is  any  compound  which  in 
the  presence  of  a  dissociating  solvent  yields  hydrogen  ions. 
Every  compound  which  dissociates  thus  is  an  acid,  and  any 
compound  which  does  not  thus  dissociate  is  not  an  acid. 

Some  of  the  consequences  of  this  definition  are  interesting. 
No  pure,  homogeneous  substance  is  appreciably  dissociated 
at  ordinary  temperatures.  Therefore,  no  pure,  homoge- 
neous substance,  in  terms  of  this  definition,  could  be  an 
acid,  at  least  of  any  appreciable  strength.  Is  this  conclusion 
hi  keeping  with  the  facts?  It  is.  As  we  have  seen,  it  has 
been  shown  that  pure,  dry  hydrochloric  acid,  even  when 
liquefied,  does  not  decompose  carbonates.  Pure,  dry, 
hydrochloric  acid  gas,  when  dissolved  in  a  nondissociating 
solvent  such  as  chloroform  or  benzene,  does  not  decompose 
carbonates,  and  does  not  even  color  blue  litmus  red. 

Sulphuric  acid,  when  properly  dried,  does  not  color  blue 
litmus  red,  and  there  are  many  other  examples  illustrating 
the  same  point.  All  of  these  examples  tend  to  confirm  the 
definition  of  an  acid  given  above. 

The  older  definitions  of  an  acid  as  a  compound  which 
would  neutralize  a  base,  and  then,  of  a  base  as  a  compound 


AQUEOUS  SOLUTIONS  OF  ACIDS,  BASES,  AND  SALTS    137 

which  would  neutralize  an  acid,  could  scarcely  be  regarded 
as  definitions  hi  the  accepted  use  of  that  term. 

Logically  more  correct,  but  from  the  standpoint  of  fact 
entirely  unreliable,  is  the  definition  of  an  acid  as  a  compound 
whose  hydrogen  could  be  replaced  directly  by  metals.  When 
ammonia  gas  is  passed  over  metallic  sodium,  one  hydrogen 
of  the  ammonia  is  replaced  directly  by  the  sodium,  sodamide 
being  formed  and  hydrogen  liberated  hi  the  sense  of  the 
following  equation, 


Na=NaNH 


and,  notwithstanding  some  work  which  would  indicate  acid 
properties  in  aqueous  solutions  of  ammonia,  we  would  hesitate* 
to  speak  of  dry  ammonia  gas  as  an  acid. 

All  of  these  older  definitions  of  an  acid  are  entirely  inade- 
quate, in  that  they  take  into  account  what  an  acid  usually 
does,  and  not  at  all  what  it  is. 

What  is  a  Base?  —  We  define  a  base  as  a  compound 
which,  hi  the  presence  of  a  dissociating  solvent,  yields  hy- 
droxyl  ions. 

In  terms  of  these  definitions  an  acid  would  be  repre- 
sented in  general  by  the  equation, 

RH=R  +  H 

in  which  R  may  be  any  atom  or  group  of  atoms,  usually  a 
group.  The  general  equation  for  a  base  would  be 

R1OH=R1+OH 

hi  which  RI  may  be  any  atom  or  group,  but  is  usually  a 
metal  atom. 

Action  of  Acids  on  Bases.  —  What  takes  place  when  we 
bring  an  acid  in  contact  with  a  base?  It  has  long  been 
known  that  the  one  neutralizes  the  other.  Each  destroys 
the  characteristic  properties  of  the  other.  How? 

Whenever  an  acid  acts  on  a  base  water  is  formed.  We 
can  see  at  once  how  this  would  take  place,  if  we  examine  the 


138  THE  NATURE  OF  SOLUTION 

above  equations  for  acids  in  general  and  for  bases  in  general, 

R  +  H  +  OH  +  R!  =  H20  +  R  +  Ri 

+ 

The  anion  of  the  acid  R,  and  the  cation  of  the  base  Ri, 
remain  after  neutralization  in  exactly  the  same  condition 
as  before,  i.e.,  they  remain  in  the  solution  in  the  ionic  state, 
uncombined  with  one  another.  The  hydrogen  ion  of  the 
acid  and  the  hydroxyl  ion  of  the  base  combine  and  form  a 
molecule  of  water,  and  this  is  what  takes  place  and  all  that 
takes  place  in  the  process  of  neutralization.  Such  are  the 
conclusions  from  the  dissociation  theory.  Are  they  true? 
It  is  easy  to  show  that  if  we  are  dealing  with  dilute  solutions 
there  is  no  salt  formed.  The  anion  of  the  acid  and  the 
cation  of  the  base  remain  separate,  in  a  word,  the  salt  which 
would  be  formed  if  these  ions  combined,  is  not  formed  at  all, 
but  is  completely  dissociated.  To  prove  this,  it  is  only 
necessary  to  measure  the  dissociation  of  the  solution  by  any 
of  the  well-recognized  methods.  It  is  well  known,  of  course, 
that  if  such  a  solution  is  evaporated,  and  the  water  which 
holds  the  ions  apart  removed,  they  will  combine  and  form  a 
salt.  This,  of  course,  has  nothing  to  do  with  their  condition 
in  dilute  solution. 

To  test  the  second  point,  whether  the  hydrogen  and 
hydroxyl  ions  do  combine  and  form  water,  is  not  so  simple. 
The  water  formed  in  the  process  of  neutralization  would 
be  in  the  presence  of  a  much  larger  quantity  of  water,  and 
would  be  difficult  to  detect.  The  question  raised  here  is 
really  a  fundamental  one  for  chemistry.  Do  hydrogen  and 
hydroxyl  ions,  in  general,  when  brought  into  the  presence 
of  one  another  combine,  or  do  they  remain  separate? 

Some  light  can  be  thrown  on  this  question  in  the  follow- 
ing way.  If  hydrogen  and  hydroxyl  ions  combine,  water  is 
formed  and  the  question  resolves  itself  into  this,  is  pure 
water  dissociated?  This  question  has  been  answered  by 
determining  the  conductivity  of  pure  water,  and  it  has  been 
found  that  pure  water  is  only  very  slightly  dissociated. 


AQUEOUS  SOLUTIONS  OF  ACIDS,   BASES,  AND  SALTS   139 

Indeed,  its  dissociation  is  so  slight  that,  under  ordinary 
conditions,  it  is  negligible.  This  fact,  as  far  as  it  goes, 
answers  the  question  thus:  hydrogen  and  hydroxyl  ions 
cannot  remain  in  the  presence  of  one  another  to  any  appre- 
ciable extent,  uncombined.  There  are  six  or  eight  other 
independent  lines  of  evidence  bearing  on  this  same  point; 
and  they  all  lead  to  the  same  conclusion,  that  hydrogen 
ions  combine  with  hydroxyl  whenever  the  two  are  brought 
together.  It  would  lead  us  too  far  to  take  these  up  here 
in  any  detail. 

Importance  of  This  Fact  for  Chemistry.  —  The  im- 
portance of  the  above  fact  for  chemistry  it  is  difficult  to 
overestimate.  Take  first  the  reaction  just  discussed.  If  hy- 
drogen ions  did  not  combine  with  hydroxyl  ions,  then  an 
acid  would  not  neutralize  a  base.  In  the  formation  of  a  salt 
the  neutralization  of  an  acid  by  a  base  is  the  first  step.  It 
is  a  necessary,  but  not  a  sufficient  condition.  To  get  the 
salt  we  must  concentrate  the  solution  to  let  the  ions  of  the 
salt  which  are  present  combine.  The  first  step  in  salt 
formation  from  acids  and  bases,  however,  is  the  union  of  the 
hydrogen  ion  of  the  acid  with  the  hydroxyl  ion  of  the  base; 
and  we  all  know  the  importance  of  salts  for  chemistry. 

Indeed,  if  we  glance  over  chemical  reactions  in  general, 
we  may  be  surprised  to  find  in  how  many  of  them  water  is 
formed;  and  it  is  usually  formed  from  the  hydrogen  of  one 
substance  combining  with  the  hydroxyl  of  the  other.  It  is 
this  very  union  of  hydrogen  with  hydroxyl  which  often 
causes  the  reaction  hi  question  to  take  place.  It  is  not  too 
much  to  state  that  a  large  percentage  of  all  the  reactions 
known  to  the  chemist,  owe  their  existence  to  the  fact  that 
hydrogen  and  hydroxyl  ions  cannot  remain  uncombined  hi 
appreciable  quantity  hi  the  presence  of  one  another. 

Heat  Evolved  When  Acids  and  Bases  React.  —  The  dis- 
sociation theory  applied  to  the  neutralization  of  acids  by  bases 
says,  that  all  that  takes  place  is  the  union  of  the  hydrogen 
ion  of  the  acid  with  the  hydroxyl  ion  of  the  base,  forming 
water.  This  can  readily  be  tested  in  the  following  way. 


140  THE  NATURE  OF  SOLUTION 

If  the  above  conclusion  is  correct,  then  every  process  of 
neutralization  is  just  like  every  other  process.  Neutraliza- 
tion is  one  and  the  same  act,  regardless  of  the  nature  of  the 
acid  and  regardless  of  the  nature  of  the  base  —  consisting 
simply  hi  the  formation  of  a  molecule  of  water.  If  this  is 
true,  what  quantities  of  heat  are  set  free  when  we  neu- 
tralize equivalent  quantities  of  different  acids  with  dif- 
ferent bases?  The  conclusion  is  obvious.  The  heat  set 
free  by  the  neutralization  of  equivalent  quantities  of  dilute 
solutions  of  strong  acids  with  strong  bases,  must  be  constant. 
The  solutions  must  be  dilute,  and  the  acids  and  bases  both 
strong  in  order  that  the  dissociation  may  be  complete. 

This  conclusion  would  seem  to  ask  a  good  deal  of  any 
theory.  That  dilute  solutions  of  all  strong  acids  and  all 
strong  bases  should,  when  brought  together,  liberate  exactly 
the  same  amount  of  heat  seemed  highly  improbable  until  it 
was  tested  experimentally.  What  are  the  facts?1 

Heat 

Nitric  acid  +  sodium  hydroxide  13,680  cal.2 

Chloric  acid  +  sodium  hydroxide  13,760    " 

Hydrochloric  acid  +  sodium  hydroxide  13,740    " 

Formic  acid  +  sodium  hydroxide  13,400    ' 

When  the  base  is  kept  constant  and  the  acid  varied,  the 
heat  of  neutralization  is  a  constant  to  within  the  limit  of 
error  of  thermochemical  methods. 

Let  us  now  take  the  next  step,  keep  the  acid  constant 
and  vary  the  nature  of  the  base  brought  in  contact  with  it. 

Sodium  hydroxide  and  hydrochloric  acid         13,740  cal. 
Calcium  hydroxide  and  hydrochloric  acid        13,950    ' 
Strontium  hydroxide  and  hydrochloric  acid     13,800    " 

These  results  suffice  to  show  that  we  can  vary  the  acid 
or  vary  the  base  as  we  like,  provided  we  keep  within  the 
category  of  strong  acids  and  strong  bases;  the  heat  of  neu- 
tralization is  a  constant.  This  constant  is  about  13,700 
calories,  which  is  of  course  the  heat  of  combination  of  the 

1  Ostwald:  Lehrb.  d.  allg.  Chem.,  vol.  2,  I. 

2  A  "calorie"  is  the  amount  of  heat  required  to  raise  the  temperature  of 
one  gram  of  water  one  degree  centigrade. 


AQUEOUS  SOLUTIONS  OF  ACIDS,   BASES,   AND  SALTS    141 

hydrogen  with  the  hydroxyl  ions.  The  conclusion  from 
the  dissociation  theory  is  thus  completely  substantiated  by 
the  experimental  facts. 

Exceptions  Presented  by  Weak  Acids  and  Weak  Bases. 

—  It  was  stated  above  that  the  law  of  the  constant  heat  of 
neutralization  holds  only  for  strong  acids  and  strong  bases. 
The  reason  for  this  is  almost  obvious.    If  the  acid  is  weak  it 
would  not  be  completely  dissociated  in  the  solution.    If  the 
base  is  weak  it  also  would  not  be  completely  dissociated. 
If  the  solutions  of  acid  and  base  brought  together  are  not 
completely  dissociated,  as  the  dissociated  portions  neutralize 
one  another  the  undissociated  portions  will  dissociate.     But 
the  dissociation  of  molecules  into  ions  has  a  thermal  value, 

—  heat  is  either  liberated  or  absorbed.    When  such  incom- 
pletely dissociated  solutions  are  brought  together,  we  have 
set  free  not  only  the  heat  of  neutralization,  which,  as  we  have 
seen,  is  the  heat  of  combination  of  hydrogen  and  hydroxyl 
ions,  but  in  addition  the  heat  of  ionization  which  may  be 
either  a  positive  or  a  negative  quantity. 

When  weak  acids  are  brought  hi  contact  with  weak  bases, 
we  should  therefore  expect  not  a  constant  heat  of  neutraliza- 
tion, but  a  value  which  might  be  either  greater  or  less  than 
the  thermal  constant  for  strong  acids  and  strong  bases,  and 
which  would  vary  somewhat  for  every  weak  acid  and  for 
every  weak  base. 

The  facts  are: 

Acetic  acid  and  ammonium  hydroxide  11.900  cal. 

Acetic  acid  and  barium  hydroxide  13.400   " 

Propipnic  acid  and  barium  hydroxide  13.400  " 

Valeric  acid  and  sodium  hydroxide  14.000   " 

Here  again,  the  conclusion  from  the  theory  of  electrolytic 
dissociation  is  fully  confirmed  by  the  facts  of  experiment. 

Theory  of  Electrolytic  Dissociation  as  a  Correlator  of 
Facts.  —  The  above  relations  suggest  another  use  of  a  law 
or  generalization  such  as  the  theory  of  electrolytic  disso- 
ciation. Before  this  theory  told  us  what  was  meant  by 
neutralization,  we  had  as  many  separate  problems  of  neutrali- 


142  THE  NATURE  OF  SOLUTION 

zation  to  deal  with  as  we  had  individual  acids  and  individual 
bases.  Every  act  of  neutralization  was  different  from  every 
other  act,  because  a  different  salt  was  formed  whenever 
we  varied  the  acid  and  whenever  we  varied  the  base. 

Now  such  is  not  the  case.  All  acts  of  neutralization  are 
one  act.  What  takes  place  in  every  case,  and  all  that  takes 
place  if  the  solutions  are  dilute,  is  the  formation  of  a  mole- 
cule of  water.  No  salt  is  formed  unless  the  solutions  con- 
taining the  ions  which  would  form  the  salt  is  evaporated  or 
otherwise  concentrated. 

Thus,  all  processes  of  neutralization,  and  there  are  as 
many  as  there  are  acids  and  bases,  become  one  process,  and 
the  whole  problem  of  neutralization  becomes  one  problem, 
which  immensely  simplifies  the  whole  subject. 

This  coordination  and  correlation  of  heterogeneous  and 
often  chaotic  and  meaningless  facts  is  one  of  the  most 
important  functions  of  a  generalization  or  law.  A  law,  then, 
tends  not  only  to  convert  empiricism  and  system  into  science, 
but  also,  by  correlation,  to  simplify  greatly  the  branch  of 
science  to  which  it  applies. 

Thermoneutrality  of  Solutions  of  Salts.  —  While  discuss- 
ing the  thermal  changes  that  take  place  in  solution,  one 
other  property  of  solutions  should  be  mentioned  in  this  con- 
^nection.  Very  dilute  solutions  of  neutral  salts  when  mixed 
show  no  thermal  change.  This  fact,  when  first  discovered, 
was  very  perplexing.  Every  chemical  change,  it  was  said, 
was  accompanied  by  a  thermal  change.  The  fact  is  correct, 
but  we  shall  see  that  the  statement  of  it  was  faulty,  in  that 
it  confused  cause  and  effect.  Here  was  certainly  a  chemical 
act  and  yet  no  thermal  change,  —  a  chemical  act,  because 
if  we  mix  solutions  of  two  salts,  say  calcium  chloride  and 
sodium  nitrate,  and  evaporate  the  mixture  to  dryness  we 
obtain  four  salts  —  calcium  chloride,  calcium  nitrate,  sodium 
chloride  and  sodium  nitrate.  If  we  start  with  two  salts  and 
end  with  four,  there  must  have  been  a  chemical  act  some- 
where in  the  process. 

This  relation  was  discovered  long  before  we  had  the 


AQUEOUS  SOLUTIONS  OF  ACIDS,   BASES,  AND  SALTS    143 

theory  of  electrolytic  dissociation  to  explain  it.  The  explana- 
tion now  is  not  only  simple,  but  this  law  is  a  necessary  con- 
sequence of  the  theory  of  electrolytic  dissociation.  Calcium 
chloride  in  solution  is  dissociated  thus, 

CaCl2  =  C&,  Cl,  Cl 
sodium  nitrate  in  solution  thus, 

NaN03  =  Na,  NO3 

When  we  mix  dilute  solutions  of  these  two  salts,  the  ions 
of  both  remain  in  the  mixture  in  exactly  the  same  condition 
of  freedom  as  in  the  separate  solutions.  There  is  no  com- 
bination—  no  salt  is  formed,  no  chemical  act  takes  place,  and 
there  should  be  no  thermal  change. 

When  the  solution  is  evaporated  and  the  water  which 
held  the  ions  apart  removed,  they  combine  in  all  possible 
ways  and  form  the  four  salts.  This,  however,  has  nothing 
to  do  with  what  takes  place  on  simply  mixing  the  dilute 
solutions  of  the  two  salts.  Thus,  the  law  of  the  ther- 
moneutrality  of  salts  presents  no  further  difficulties,  now 
that  we  have  the  theory  of  electrolytic  dissociation. 

Importance  of  Energy  Changes  for  Chemistry.  —  We 
have  just  discussed  the  energy  change  which  takes  place 
when  acids  are  neutralized  by  bases,  and  have  pointed  out 
that  every  chemical  reaction  is  always  accompanied  by  a 
thermal  change,  as  was  stated.  This  raises  the  question, 
of  what  significance  are  the  energy  changes  for  chemistry? 

When  the  law  of  the  conservation  of  energy  was  dis- 
covered, it  was  supposed  to  be  simply  a  law  of  physics,  and 
chemists  paid  comparatively  little  attention  to  it.  This  was 
due  primarily  to  the  fact  that  chemists  of  that  period  did 
not  pay  very  serious  attention  to  physics  in  general,  being 
more  interested  hi  the  preparation  of  compounds,  and  hi  the 
working  out  of  their  composition  and  constitution,  than  in 
discovering  the  physical  principles  that  underlie  chemical 
science.  This  was  perfectly  natural  when  we  consider  the 
development  of  chemistry  at  the  tune  to  which  we  are  refer- 


144  THE  NATURE  OF  SOLUTION 

ring.  It  was  not  until  the  great  French  chemist,  Berthelot, 
pointed  out  the  importance  of  energy  changes  for  chemistry, 
and  made  his  famous  thermochemical  measurements  from 
which  he  deduced  fundamental  laws,  that  chemists  began  to 
see  the  significance  of  changes  hi  energy  for  the  whole  science 
of  chemistry. 

Since  the  tune  of  Berthelot's  thermochemical  work,1 
we  have  recognized  that  the  cause  of  all  chemical  reaction 
is  to  be  found  hi  the  different  amounts  and  potentials  of 
intrinsic  energy  in  the  substances  brought  together.  When 
this  difference  is  sufficiently  great,  some  of  the  intrinsic 
energy  is  converted  into  heat,  light,  or  electricity  —  always 
some  into  heat,  and  the  substances  rearrange  themselves 
in  new  combinations  which  are  more  stable  under  the  new 
conditions.  A  chemical  reaction  is,  then,  not  "  accom- 
panied" by  a  thermal  change,  but  is  caused  by  it.  We 
could  more  truly  say  that  the  thermal  change  is  "  accom- 
panied" by  the  chemical  reaction. 

As  soon  as  the  importance  of  energy  changes  for  chemistry 
was  pointed  out,  chemists  began  to  realize  that  the  law  of 
the  conservation  of  energy  is  as  important  for  chemistry  as 
it  is  for  physics.  The  two  laws,  the  law  of  the  conservation 
of  mass  and  the  law  of  the  conservation  of  energy,  lie  right 
at  the  foundation  of  chemical  science;  chemistry  being  quite 
as  much  dependent  for  its  existence  upon  energy  changes  as 
is  physics. 

Color  of  Dissolved  Substances.  —  Many  solutions  in 
water  are  colorless,  and  others  are  colored.  Why  this  dif- 
ference? Light  is  a  series  of  vibrations  of  different  wave 
lengths  hi  the  hypothetical  ether.  White  light  contains  a 
large  number  of  such  vibrations.  When  all  of  the  ether 
vibrations  pass  on  through  the  solution,  it  is  white  or  col- 
orless as  we  say.  When  all  of  these  vibrations  are  cut  off 
or  stopped,  the  solution  is  opaque  or  black.  When  some  of 
the  vibrations  are  cut  off  and  others  allowed  to  pass  on 
through,  the  solution  is  colored;  the  color  depending  upon 

1  Essai  de  Mechanique  Chimique. 


AQUEOUS  SOLUTIONS  OF  ACIDS,   BASES,  AND  SALTS    145 

the  particular  wave-lengths  of  the  vibrations  that  are  trans- 
mitted. If  the  shorter  wave-lengths  are  stopped,  the  solu- 
tion has  the  color  of  the  longer  wave-lengths  which  pass  on 
through.  If  the  longer  wave-lengths  are  cut  out,  the  solu- 
tion has  the  color  of  the  shorter  wave  lengths  that  are  trans- 
mitted. 

The  solution  may  cut  out  some  of  the  longer  and  some 
of  the  shorter  wave-lengths,  in  which  case  the  solution  would 
have  the  color  of  the  transmitted  light,  whatever,  it  is.  It 
is  thus  obvious  that  solutions  may  have  any  color. 

Absorption  of  Light  due  to  Resonance.  —  The  question 
arises,  what  is  meant  by  absorption  of  light?  To  use  the 
prevailing  theory,  how  can  vibrations  in  the  ether  be  ab- 
sorbed or  stopped? 

If  the  ether  vibrations  find  something  hi  the  solution 
which  they  can  set  vibrating  with  the  same  period  as  their 
own  —  can  throw  into  resonance  with  themselves  —  their 
energy  is  expended  in  setting  this  something  vibrating,  and 
the  original  vibration  is  stopped.  Those  vibrations  hi  the 
ether  which  do  not  find  hi  the  solution  anything  which  they 
can  set  vibrating  with  their  own  period,  are  not  stopped,  but 
pass  on  through.  Resonance  produces  opacity;  the  lack  of 
resonance,  transparency. 

Color  in  Solution  May  be  Ionic  or  May  be  Molecular.  — 
The  color  of  solutions  acquired  a  new  interest  after  the  theory 
of  electrolytic  dissociation  was  proposed.  If  dilute  solutions 
of  electrolytes  contain  only  ions,  then  all  the  properties  of 
such  solutions,  including  their  color,  must  be  due  to  the  ions. 
This  could  readily  be  tested,  and  it  was  done  by  Ostwald1  in 
the  following  manner. 

Select  a  series  of  salts  of  an  acid  with  a  colored  anion,  say 
fluorescein,  the  salts  having  colorless  cations;  and  see  whether 
the  solutions  of  these  different  salts  all  have  the  same  color. 
This,  of  course,  could  not  be  tested  by  the  eye,  since  the 
eye  as  a  measure  of  color  is  most  deceptive.  A  spectroscope 
must  be  used  for  this  purpose,  and  Ostwald  used  a  prism. 

1  Zeit.  phys.  Chem.,  9,  579  (1892). 


146  THE  NATURE  OF  SOLUTION 

The  absorption  bands  for  the  different  solutions  were 
photographed  the  one  above  the  other.  These  bands  cor- 
respond to  the  wave-lengths  of  light  that  were  cut  out  or 
absorbed  by  the  solutions. 

Take  for  example  the  salts  of  fluorescein.1  The  ammo- 
nium, cobalt,  magnesium,  cadmium,  barium,  manganese, 
lithium,  nickel,  potassium,  and  zinc  salts  were  studied. 

SALTS  OF  FLUORESCEIN 

X  =  A.U. 

Ammonium  2866 

Cobalt  2867 

Magnesium  2865 

Cadmium  2866 

Barium  2866 

Manganese  2867 

Lithium  2865 

Nickel  2866 

Potassium  2866 

Zinc  2866 

The  absorption  bands  for  all  of  these  salts  of  fluorescein 
coincide  almost  to  within  the  limit  of  experimental  error. 
Ostwald  carried  out  similar  measurements  with  salts  of 
eosin,  iodoeosin,  and  other  acids  with  colored  anions,  the 
salts  chosen  having  colorless  cations.  Similar  results  were 
obtained  hi  every  case;  the  absorption  of  the  different 
salts  of  any  one  acid  being  the  same  to  within  the  limit  of 
experimental  error. 

Having  studied  the  salts  of  acids  with  colored  anions,  the 
cations  being  colorless,  Ostwald  turned  to  salts  of  bases 
with  colored  cations,  the  anions  being  colorless.  Thus, 
he  studied  the  p-rosaniline  salts  and  aniline  violet  salts 
of  twenty  acids  with  colorless  anions.  He  photographed 
the  results  for  each  base,  the  one  above  the  other,  and  found 
that  the  absorption  lines  coincided  to  within  experimental 
error.  A  few  of  his  results  for  aniline  violet  with  a  number  of 
acids  are  given  hi  the  table  on  the  opposite  page.2 

1  Zeit.  phys.  Chem.,  9,  587  (1892).  2  Ibid.,  599. 


AQUEOUS  SOLUTIONS  OF  ACIDS,  BASES,  AND  SALTS  147 

SALTS  OF  ANILINE  VIOLET 

X  =  A.U. 

Acetate  2534 

Benzoate  2534 

Hydrochloride  2533 

Nitrate  2534 

Butyrate  2533 

Trichloracetate  2533 

Glycolate  2534 

Lactate  2533 

These  results  obtained  by  Ostwald  fall  right  in  line  with 
the  dissociation  theory  for  electrolytes.  The  colors  of  these 
substances  are  the  colors  of  the  ions. 

We  know,  however,  that  molecules  also  can  have  color. 
Many  solutions  of  non-electrolytes  which  are  completely 
undissociated  are  colored.  Again,  many  substances  hi  the 
pure,  dry  condition  are  colored,  and  these  are,  of  course, 
undissociated. 

In  the  present  connection  it  should  be  pointed  out  that 
the  crystals  of  many  substances  which  crystallize  with  water 
are  colored,  and  the  crystals  have  approximately  the  same 
color  as  their  aqueous  solutions.  What  is  the  explanation? 
The  molecules  are  dissociated  in  their  water  of  crystalliza- 
tion, and  the  ions  thus  formed  show  their  characteristic 
color. 

Cause  of  Color  in  Solution.  —  We  have  seen  that  color 
is  due  to  resonance.  Those  wave-lengths  of  light  which  can 
find  something  in  the  solution  which  they  can  throw  into 
resonance,  are  cut  out  or  absorbed,  the  remaining  vibrations 
passing  through  and  giving  the  characteristic  color  to  the 
solution.  This  raises  the  question,  what  is  it  in  solution  that 
is  thrown  into  resonance  by  the  light?  Is  it  the  molecule, 
the  ion,  the  atom,  or  the  electron  or  electrical  charge1  of 
which  all  the  atoms  are  made? 

We  shall  see  in  the  last  chapters  of  this  book  that  work 
in  this  laboratory  has  made  it  almost  certain  that  absorption 
hi  solution  is  not  due  directly  to  the  molecules,  nor  to  the 
atoms,  and  not  even  to  the  ions,  but  is  due  to  the  electrons. 

1  See  Author's  Electrical  Nature  of  Matter  and  Radioactivity,  3d  edition 
(The  Van  Nostrand  Co.). 


148  THE  NATURE  OF  SOLUTION 

It  is  interesting  to  note  that  this  was  predicted  by  Ost- 
wald  long  before  Thomson  had  done  his  epoch-making  work 
on  the  electron.  Ostwald  pointed  out  the  importance  of  the 
electrical  charge  on  the  ion  as  conditioning  its  color,  and  this 
electrical  charge  we  now  know  is  nothing  but  the  electron  of 
Thomson. 

Take  as  an  illustration  of  this  point  the  well-known  salts, 
potassium  ferrocyanide,  K4Fe(CN)6,  and  potassium  ferricy- 
anide,  K3Fe(CN)6.  Ostwald  supposed  that  these  compounds 
dissociated  as  follows: 

K4  Fe(CN)6  =  K,K,K,K,  Fe(CN)6 

and  K3Fe(CN)6  =K,K,K,  Fe(CN)6 

The  color  of  the  ferrocyanide  is  yellow;  while  the  color 
of  the  ferricyanide  is  greenish;  and  the  latter  has  many  times 
the  coloring  power  of  the  former.  Why  this  difference? 

The  color  of  both  solutions  is  due  to  the  ferrocyanogen 
ions  since  the  potassium  ion  is  colorless.  The  ferrocyanogen 
ion  from  the  ferrocyanide  has  four  negative  charges  upon  it, 
while  the  ferrocyanogen  ion  from  the  ferricyanide  has  three 
negative  charges  upon  it.  Ostwald  would  account  for  the 
difference  in  the  color  as  due  to  this  difference  in  the  num- 
ber of  charges  which  these  ions  carry;  and  this,  as  has  been 
stated,  is  in  perfect  accord  with  our  own  recent  work  on  the 
absorption  spectra  of  solutions. 

Color  Changes  and  Volumetric  Analysis  —  Indicators.  — 
The  whole  science  of  volumetric  analysis  is  based  upon 
certain  changes  in  color  of  certain  substances  known  as 
indicators.  These  substances  have  been  used,  and  these 
changes  in  color  employed,  in  volumetric  analysis  from  the 
early  history  of  this  important  branch  of  quantitative 
chemistry.  The  indicators  were,  however,  used  for  a  long 
time  largely  mechanically.  Certain  things  were  done  and 
certain  results  obtained  without  knowing  just  why. 

It  remained  for  Ostwald  to  clear  up  this  subject  as  he  has 
cleared  up  so  many  others.  The  Ostwald  theory  of  indica- 


AQUEOUS  SOLUTIONS  OF  ACIDS,  BASES,  AND  SALTS    149 

tors  has  taken  this  subject  from  the  rank  of  pure  empiricism 
and  placed  it  for  the  first  time  upon  a  really  scientific  basis. 
Let  us  see  what  it  is. 

The  Ostwald  theory,  of  course,  is  based  upon  the  theory 
of  electrolytic  dissociation  as  a  starting  point.  Indicators 
are  either  weak  acids  or  weak  bases.  Further,  the  dis- 
sociated compound,  the  ions,  must  have  a  different  color 
from  the  undissociated  molecules.  These  are  the  keys  to 
the  whole  theory. 

Take  phenolphthalein;  its  molecules  are  colorless  and  its 
anion  colored.  Phenolphthalein,  being  a  weak  acid,  in 
the  presence  of  a  strong  acid,  or  in  the  presence  of  water 
alone,  is  undissociated.  When  a  strong  base  is  added  to 
phenolphthalein  it  dissociates,  yielding  the  anion  with  its 
characteristic  color.  This  compound  illustrates  the  type 
of  indicator  where  the  molecule  is  colorless  and  the  anion 
colored. 

Litmus  illustrates  another  type  of  indicator.  Like 
phenolphthalein,  litmus  is  a  weak  organic  acid;  but  unlike 
phenolphthalein  the  molecules  of  litmus  acid  are  colored, 
and  are  colored  red.  When  a  base  is  added  to  litmus  acid, 
the  molecules  are  caused  to  dissociate  and  set  free  the  anion 
of  the  acid  which  is  colored  blue  —  the  molecules  have  one 
color  and  the  ions  another. 

Another  acid  indicator  which,  however,  is  a  much  stronger 
acid  than  litmus,  is  methyl  orange.  The  molecules  are  red, 
and  therefore  this  indicator  in  the  presence  of  a  strong  acid 
is  red.  The  anions  are  yellow.  Therefore,  hi  the  presence  of 
a  base  this  indicator  dissociates,  liberating  the  yellow  anions. 

Exactly  the  same  relation  obtains  for  methyl  red,  which  is 
very  closely  allied  chemically  to  methyl  orange.  Here  the 
molecules  are  red,  and  the  reaction  is  therefore  red  in  the 
presence  of  acids.  The  anion  is  yellow,  and  this  indicator 
therefore  reacts  yellow  in  the  presence  of  a  base. 

Another  acid  indicator  is  corallin,  or  rosolic  acid.  It  is  a 
weak  organic  acid,  the  molecules  being  yellow.  When  a 
base  is  added  the  red  anion  is  set  free. 


150  THE  NATURE  OF  SOLUTION 

Stieglitz1  s  Views  in  Reference  to  Phenolphthakin.  —  The 
view  of  Stieglitz1  in  reference  to  the  action  of  phenolphthal- 
ein  as  an  indicator  goes  further  than  the  theory  of  Ostwald. 
The  latter,  from  what  has  been  said  above,  simply  assumes 
that  the  weak,  colorless  phenolphthalein  molecules  disso- 
ciate into  hydrogen  ions  which  combine  with  the  hydroxyl 
ions  of  the  base,  and  into  an  anion  which  has  the  color 
characteristic  of  this  indicator  hi  the  presence  of  bases. 

Stieglitz,  commenting  on  the  Ostwald  theory  of  indicators 
says,  "It  is  extremely  probable,  moreover,  that  this  theory  is 
wrong  hi  so  far  as  the  interpretation  of  the  cause  of  the 
change  of  color  is  concerned."  "While  most  likely  wrong 
in  regard  to  the  one  question  of  change  of  color,  Ostwald, 
hi  his  theory  of  indicators,  has  undoubtedly  laid  down 
correctly  the  guiding  principles  of  the  proper  theoretical 
treatment  of  the  second,  scientifically  far  more  important, 
question  concerning  the  varying  sensitiveness  of  the  in- 
dicators to  acids  and  bases." 

"In2  view  of  what  chemists  have  known  for  over  a 
quarter  of  a  century  about  the  intimate  connection  between 
color  production  and  constitution  of  organic  compounds, 
the  explanation  that  phenolphthalein  without  a  single  chro- 
mophoric  group  should  become  intensely  red  by  forming  the 
ion  without  any  chromophoric  group  appeared  from  the 
outset  as  extremely  unlikely." 

Stieglitz  then  points  out  that  it  had  been  shown  by 
Bernthsen  that  when  phenolphthalein  forms  a  salt,  as  when 
treated  with  an  alkali,  there  is  an  internal  rearrangement, 
giving  rise  to  the  highly  colored  quinoid  (:C6H4:0)  group, 
which  explains  the  appearance  of  the  color.  "The  sodium 
salt  is,  no  doubt,  incidentally  ionized  in  solution,  but  that 
this  ionization  is  merely  a  coincidence  and  not  a  cause  is 
established  by  the  fact  that  the  solid,  dry,  non-ionized  silver 
salt  is  also  intensely  colored  (violet)."  Says  Stieglitz,  "more 
recently  Hantzsch  has  shown  that,  in  general,  the  change  of 
colorless  organic  compounds  into  highly  colored  salts  is 

1  Journ.  Amer.  Chem.  Soc.,  26,  1112  (1903.)  2  Ibid.,  1114. 


AQUEOUS  SOLUTIONS  OF  ACIDS,   BASES,  AND  SALTS    151 

invariably  accompanied  by  a  modification  of  the  constitution 
affecting  a  chromophoric  group." 

According  to  Stieglitz  the  ionization  alone  is  not  the 
cause  of  the  color,  but  the  ionization  and  the  internal  change 
in  constitution  take  place  simultaneously,  and  it  is  the  latter 
which  is  the  cause  of  the  color. 

In  the  use  of  indicators  hi  quantitative  analysis  it  is 
important  to  know  the  degree  of  dissociation  of  the  weak 
acids  and  weak  bases  used  as  indicators.  The  so-called 
"constants"  of  indicators  are  a  function  of  their  dissocia- 
tions. These  constants1  have  recently  been  worked  out 
radiometrically  in  this  laboratory  for  phenolphthalein, 
methyl  orange,  and  corallin. 

Amphoteiic  Compounds  —  Importance  in  Chemistry.  — 
We  have  considered  acids  as  dissociating  always  into  a  hy- 
drogen cation,  and  into  an  anion  whose  composition  depends 
on  the  particular  acid  hi  question.  Wherever  we  have  acid 
properties  we  have  hydrogen  ions,  and  wherever  we  have 
hydrogen  ions  we  have  acid  properties. 

Similarly,  bases  have  been  referred  to  as  compounds 
which,  hi  the  presence  of  a  dissociating  solvent,  yield 
hydroxyl  ions.  Basicity  and  hydroxyl  ions  always  go  hand 
in  hand. 

It  might  be  inferred  from  this  that  a  compound  which 
under  one  set  of  conditions  yields  hydrogen  ions,  under  all 
conditions  would  dissociate  yielding  ions  of  hydrogen;  and, 
similarly,  a  compound  which  under  one  set  of  conditions 
would  split  off  hydroxyl  ions,  would  under  all  conditions 
yield  the  hydroxyl  ion.  Such  is  not  the  case.  A  compound 
may  under  some  conditions  dissociate  into  hydrogen  ions, 
while  under  other  conditions  the  same  compound  may 
dissociate  yielding  hydroxyl  ions.  Such  a  compound  which 
is  capable  of  dissociating  now  as  an  acid,  and  now  as  a  base, 
is  known  as  an  amphoteric  electrolyte. 

There  is  a  fairly  large  number  of  examples  of  these 
substances,  both  in  general  chemistry,  and  among  the 

1  Journ.  Amer.  Chem.  Soc.,  37,  776  and  1694  (1915.) 


152  THE  NATURE  OF  SOLUTION 

compounds  of  carbon.  We  need  only  mention  zinc  hy- 
droxide, which  in  the  presence  of  an  acid  or  hydrogen 
ions  dissociates  thus, 

Zn(OH)2=Zn,  OH,  OH 

while  in  the  presence  of  a  base  or  hydroxyl  ions,  it  disso- 
ciates thus, 

Zn(OH)2  =   ZnO2,  H,  H 

In  the  first  case  zinc  hydroxide  dissociates  as  a  base;  hi 
the  second  as  an  acid.  As  a  base  it  neutralizes  acids  and 
forms  zinc  salts;  as  an  acid  it  neutralizes  bases  and  forms 
zincates.  It  is  a  well-known  fact  that  zinc  hydroxide  dis- 
solves in  an  excess  of  sodium  hydroxide  almost  as  readily  as 
in  hydrochloric  acid. 

Take  another  example  illustrating  the  same  point  — 
lead  hydroxide.  In  the  presence  of  an  acid  it  dissociates  as 
a  base,  thus, 

Pb(OH)2  =  Pb,  OH,  OH 
In  the  presence  of  a  base,  it  dissociates  as  an  acid,  thus, 

Pb(OH)2  =  Pb=02,  H,  H 

As  a  base  lead  hydroxide  reacts  with  acids  forming  salts  of 
lead.  As  an  acid  it  reacts  with  bases  forming  plumbates. 

There  are  many  other  examples  of  amphoteric  substances 
in  general1  as  in  organic  chemistry,  but  the  above  suffice  to 
illustrate  the  principle. 

Amphoterism  and  Biological  Processes.  —  The  number 
of  amphoteric  substances  among  the  compounds  of  carbon 
is  very  large  indeed.  We  need  only  mention  asparagine, 
sarcosine,  leucine,  tyrocine,  and  especially  the  proteids  in 
general;  and  the  role  of  the  latter  in  living  processes  is  very 
important.  They  contain  both  the  carboxyl  (COOH)  and 
the  amino  (NH2)  groups,  and  can  therefore  dissociate  either 

1  Elements  of  Inorganic  Chemistry,  H.  C.  Jones,  3d  edition,  p.  389.  (The 
Macmillan  Co.) 


AQUEOUS  SOLUTIONS  OF  ACIDS,  BASES,   AND  SALTS    153 

as  acids  or  as  bases.  In  this  connection  see  L.  J.  Henderson, 
"The  Fitness  of  the  Environment";  and  S.  L.  Sorensen, 
Ergebnisse  der  Physiologie,  vol.  12  (1912). 

Hydrolysis  at  Ordinary  and  at  High  Temperatures. — 
The  only  action  of  water  on  salts  that  we  have  thus  far  dis- 
cussed is  that  of  electrolytic  dissociation.  The  molecule 
of  the  salt  is  broken  down  directly  by  water  into  a  cation  or 
cations  charged  positively,  and  into  an  anion  or  anions 
charged  negatively.  There  is  another  kind  of  dissociation 
of  salts  effected  by  water  which  must  be  discussed  here  — 
hydrolytic  dissociation  —  which  differs  fundamentally  from 
electrolytic  dissociation,  although  the  products  of  hydrolytic 
dissociation  are  always  more  or  less  electrolytically  disso- 
ciated. 

Solutions  of  certain  salts  react  acid,  e.g.,  aluminium  chlo- 
ride; while  solutions  of  sodium  carbonate  react  alkaline. 
This  fact  was  known  a  long  while  before  it  could  be  explained. 
It  was  said  that  the  former  reacts  acid  because  hydrochloric 
acid  is  a  stronger  acid  than  aluminium  hydroxide  is  a  base. 
A  residue  of  the  acid  properties  still  clings  to  the  salt.  Simi- 
larly, sodium  carbonate  reacts  alkaline  because  sodium 
hydroxide  is  a  stronger  base  than  carbonic  acid  is  an  acid. 

In  the  light  of  what  we  know  today  about  acids  and  bases, 
this,  of  course,  explains  nothing,  and  is  hardly  more  than 
words. 

We  now  know  what  is  meant  by  hydrolysis,  and  why 
solutions  of  some  salts  react  acid  and  others  basic.  Alumi- 
nium chloride,  for  instance,  is  acted  upon  by  water,  to  a 
slight  extent  at  least, 

2A1C13  +  3H2O  =  6HC1  +  2A1(OH)3 

Aluminium  hydroxide,  being  a  weak  base,  is  only  slightly 
dissociated,  and  yields  only  relatively  few  hydroxyl  ions. 
Hydrochloric  acid,  on  the  other  hand,  is  a  strong  acid  and 
is  strongly  dissociated  into  its  ions.  The  hydrogen  ions 
predominate  over  the  hydroxyl,  whence  the  acid  reaction 
of  the  solution. 


154  THE  NATURE  OF  SOLUTION 

In  the  case  of  sodium  carbonate  exactly  the  opposite  is 
true.    The  reaction  with  water  is  represented  thus, 

NasCOs  +  2H20  =  2NaOH  +  H2C03 

Sodium  hydroxide,  being  a  strong  base,  is  largely  dissoci- 
ated and  yields  a  large  number  of  hydroxyl  ions.  Carbonic 
acid,  being  a  weak  acid,  is  only  slightly  dissociated  and  con- 
sequently yields  only  a  few  hydrogen  ions.  The  hydroxyl 
ions  predominate  hi  number  over  the  hydrogen  ions  and  give 
the  alkaline  reaction  to  the  solution. 

The  question  arises,  is  hydrolysis  a  general  phenomenon? 
Are  salts  in  general  hydrolytically  dissociated  by  water? 
The  answer  is,  yes.  Salts  of  strong  acids  with  strong  bases, 
however,  undergo  such  slight  hydrolysis  that  it  can  usually 
be  disregarded.  Salts  of  weak  acids  with  weak  bases  are 
the  most  strongly  hydrolyzed.  Salts  of  weak  acids  with 
strong  bases,  and  of  weak  bases  with  strong  acids  are  hydro- 
lyzed to  some  extent,  the  magnitude  of  the  hydrolysis 
depending  on  the  actual  strengths  of  the  acid  and  base 
which  formed  the  salt  in  question. 

It  may  be  said  in  general  that  salts  of  the  alkalies  with 
strong  acids  are  so  little  hydrolyzed  by  water  at  ordinary 
temperatures,  that  the  hydrolysis  is  negligible.  Salts  of  the 
alkaline  earths,  even  with  the  strongest  acids,  are  hydrolyzed 
to  a  slight  extent,  while  salts  of  weaker  bases,  such  as 
aluminium  hydroxide,  ferric  hydroxide,  etc.,  are,  even  at 
ordinary  temperatures,  very  appreciably  hydrolyzed. 

It  will  be  noted  that  the  products  of  hydrolytic  dissocia- 
tion are  always  more  or  less  electrolytically  dissociated  by 
water  —  hydrolytic  dissociation  is  always  accompanied  by 
more  or  less  electrolytic  dissociation. 

Effect  of  Temperature  on  Hydrolytic  Dissociation. — 
The  effect  of  rise  in  temperature  on  hydrolysis  is  a  matter  of 
interest  not  only  for  the  chemist,  but  for  the  geologist. 
This  problem  has  been  studied  by  A.  A.  Noyes1  and  his  co- 
workers.  They  used  a  closed  bomb  to  hold  the  solutions  at 

1  Zeit.  phys.  Chem.,  46,  323  (1903.) 


AQUEOUS  SOLUTIONS  OF  ACIDS,  BASES,  AND  SALTS    155 

more  elevated  temperatures,  and  calculated  the  change  in 
hydrolysis  with  rise  in  temperature  from  the  change  in  the 
conductivity.  They  worked  from  0°  up  to  306°  and  at  even 
higher  temperatures. 

We  have  seen  that  the  most  strongly  hydrolyzed  salts 
are  the  salts  of  weak  acids  and  weak  bases.  Noyes  and  his 
co-workers  studied  the  effect  of  rise  hi  temperature  on  the 
hydrolysis  of  ammonium  acetate  —  a  salt  of  a  weak  acid 
and  a  weak  base.  They  found  that  for  a  rise  in  temperature 
from  0°  to  306°,  the  hydrolysis  of  ammonium  acetate  in- 
creased about  three  hundred  times. 

The  effect  of  rise  hi  temperature  on  the  hydrolytic 
decomposition  of  salts  by  water  is  of  geological  significance 
and  importance.  Well  down  beneath  the  surface  of  the 
earth  highly  heated  water-vapor  comes  in  contact  with  the 
rocks  which  are  also  heated  to  very  high  temperatures. 
The  decomposition  of  the  silicates  and  other  rocks  by  water 
at  very  high  temperatures  is  now  intelligible  hi  the  light  of 
the  work  of  Noyes  on  the  rapid  increase  in  hydrolysis  with 
rise  in  temperature. 


CHAPTER  X 

A   FEW   ELECTRICAL   PROPERTIES    OF   AQUEOUS    SOLUTIONS 
OF    ELECTROLYTES 

THE  difference  between  electrolytes  and  non-electro- 
lytes was  clearly  pointed  out  in  an  earlier  chapter.  It 
was  shown  that  electrolytes  are  those  substances  whose 
solutions  conduct  the  current,  and  that  non-electrolytes  are 
those  whose  solutions  do  not  conduct. 

Since  solutions  of  electrolytes  conduct  electricity,  the 
particles  in  such  solutions  must  be  charged  electrically. 

The  first  to  study  scientifically  and  successfully  electrical 
conduction  in  solution  was  the  great  experimenter,  Michael 
Faraday.  He  passed  different  amounts  of  electricity  through 
the  solution  of  the  same  salt,  and  determined  the  amounts 
of  thejnetal  deposited  upon  the  cathode.  In  this  way  he 
discovered  the  first  part  of  his  well-known  law,  or  his  first 
law. 

Faraday 's  First  Law.  —  Faraday  found  that  when  double 
the  amount  of  electricity  was  passed  through  a  solution, 
exactly  double  the  amount  of  salt  was  decomposed  and 
double  the  amount  of  metal  deposited  upon  the  cathode. 
He  studied  the  decompositions  of  a  large  number  of  sub- 
stances, using  different  amounts  of  current,  and  always  found 
that  the  amount  of  decomposition  was  proportional  to  the 
amount  of  the  current  passed  through  the  solution.  He 
formulated  his  first  law  as  follows:  The  amount  of  salt 
decomposed  by  the  electric  current  is  proportional  to  the  amount 
of  electricity  passed  through  the  solution. 

This  law  is  of  importance  in  that  it  gives  us  a  ready  means 
of  measuring  quantity  of  electricity.  Knowing  the  amount 
of  decomposition  of  a  salt  like  silver  nitrate  effected  by 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  157 

passing,  say,  one  coulomb  of  electricity,  it  is  only  necessary 
to  pass  the  current  in  question  through  a  solution  of  silver 
nitrate  and  weigh  the  amount  of  silver  deposited  upon  the 
cathode,  in  order  to  know  how  much  current  was  passed 
through  the  solution.  The  silver,  copper,  and  gas  voltam- 
eter are  all  based  upon  this  first  law  of  Faraday. 

Faraday's  Second  Law.  —  Faraday  then  raised  the  ques- 
tion, is  there  any  relation  between  the  amounts  of  the 
different  metals  deposited  by  the  same  amount  of  current, 
and  if  so,  what  is  this  relation? 

To  test  this,  Faraday  passed  the  same  current  through  a 
series  of  solutions  of  salts  of  different  metals  and  collected 
the  metals  on  the  several  cathodes  and  weighed  them.  He 
found  the  following  very  remarkable  relation.  The  amounts 
of  the  different  metals  thrown  out  by  the  same  current 
are  proportional  to  the  atomic  weights  of  the  metals,  pro- 
vided all  the  metals  hi  question  have  the  same  valence. 
If  they  have  different  valences,  Faraday  found  that  the 
amounts  of  the  metals  thrown  out  by  the  same  quantity  of  elec- 
tricity are  proportional  to  the  atomic  weights  of  the  metals  di- 
vided by  their  respective  valences;  and  this  is  the  second  law 
of  Faraday. 

Of  what  significance  is  this  law?  What  have  the  valences 
of  the  different  metals  to  do  with  the  decomposition  of  their 
salts  by  the  electric  current?  These  are  questions  which  the 
second  law  of  Faraday  would  naturally  suggest.  Then* 
answer  has  told  us  what  chemical  valence  really  means. 

Second  Law  of  Faraday  and  Chemical  Valence.  —  When 
a  given  current  is  passed  through  solutions  of  salts  of  metals 
having  the  same  valence,  as  already  stated,  the  amounts  of 
the  different  metals  deposited  are  proportional  to  their 
atomic  weights.  This  means  that  the  same  numbers  of 
atoms  of  all  the  different  metals  are  deposited  by  the  same 
amount  of  electricity.  This  is  the  same  as  to  say  that  all 
atoms  having  the  same  valence  carry  exactly  the  same  amount 
of  electricity. 

The  law,  however,  goes  farther  in  its  bearing  on,  and 


158  THE  NATURE  OF  SOLUTION 

explanation  of  chemical  valence.  If  the  metals  have  dif- 
ferent valences,  the  amounts  deposited  are  proportional  to 
the  atomic  weights  divided  by  the  several  valences.  This 
shows  exactly  what  is  the  difference  between  a  univalent,  a 
bivalent,  a  trivalent,  and  an  n-valent  element.  They  carry 
different  numbers  of  unit  charges  of  electricity  —  the  num- 
ber determining  the  valence  of  the  ion.  A  univalent  ion 
carries  one  unit  charge  of  electricity  or  one  free  electron;  a 
bivalent  ion  two  unit  charges  or  two  free  electrons;  a 
trivalent  ion  three  unit  charges  or  three  free  electrons; 
an  n-valent  ion  n  unit  charges  or  n  electrons. 

Chemical  valence  is  thus  shown  to  be  connected  hi  some 
fundamental  way  with  the  number  of  free  unit  charges 
of  electricity,  or  electrons,  upon  the  ions.  This  second 
law  of  Faraday  simply  shows  that  the  one  is  proportional 
to  the  other.  It  does  not  answer  the  question  whether  the 
one  is  the  cause  of  the  other.  This  has,  however,  been 
answered;  and  since  the  question  of  chemical  valence  is 
so  fundamental  for  all  chemistry,  the  answer  will  be  given 
here,  at  the  risk  of  a  slight  digression  from  the  questions 
now  more  directly  under  discussion. 

Nature  of  Chemical  Valence.  —  There  are  few  subjects 
upon  which  more  has  been  written  than  the  subject  of 
valence,  and  few  subjects  about  which  there  has  accumu- 
lated so  much  unsatisfactory  literature.  The  reason  for 
this  is  obvious  to  any  one  who  has  read  a  few  of  the  many 
papers  on  this  subject. 

The  attempt  has  frequently  been  made  to  discuss  chemi- 
cal valence  without  any  clean-cut  definition  of  the  subject 
under  discussion.  The  result  is  just  what  would  be  expected. 
When  we  start  with  nothing  definite  we  can  hope  to  lead 
to  nothing  definite.  It  is  scarcely  hypercritical  to  state 
that  many  of  the  papers  which  have  been  written  upon  the 
subject  of  valence,  when  carefully  analyzed,  are  little  more 
than  words,  largely  for  the  reason  indicated  above. 

The  second  law  of  Faraday  has  placed  the  whole  subject 
of  chemical  valence  upon  a  definite  physical  basis,  and  has 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  159 

shown  the  proportionality  between  the  valence  of  an  ion 
and  the  number  of  charges  which  it  carries. 

It  remained  for  Ostwald  to  show  the  causal  relation 
between  these  two  phenomena  —  that  chemical  valence  is 
essentially  the  chemical  expression  of  the  number  of  free 
electrons  upon  the  ions  and  then*  attractive  and  repellent 
forces.  This  was  done  by  the  following  experiment  hi  which 
chemical  valence  was  synthesized,  if  you  please,  and  out  of 
what?  Out  of  electrical  energy;  and  done  hi  such  a  way 
that  there  could  be  no  reasonable  question  as  to  the  result. 

Experiment  Demonstrating  the  Nature  of  Chemical 
Valence.  —  A  cell  was  constructed  in  the  following  manner.1 
Into  one  of  two  beakers  was  introduced  a  solution  of 
ferrous  chloride.  A  solution  of  potassium  chloride  was  in- 
troduced into  the  second  beaker;  and  into  each  beaker, 
a  platinum  wire  serving  as  an  electrode.  The  two  elec- 
trodes were  attached  to  a  galvanometer.  The  two  beakers 
were  connected  by  a  siphon  filled  with  the  solution  of  potas- 
sium chloride.  Chlorine  gas  was  conducted  into  the  solution 
of  potassium  chloride.  When  this  was  done  the  galvanometer 
showed  a  current  flowing  on  the  outside  of  the  cell  from  the 
electrode  surrounded  by  the  potassium  chloride  into  which 
chlorine  gas  has  been  conducted  to  the  electrode  surrounded 
by  the  neutral  ferrous  chloride.  The  ferrous  chloride 
quickly  showed  the  presence  of  ferric  chloride. 

The  action  of  this  cell  is  as  follows.  When  chlorine 
gas  is  conducted  into  the  solution  of  potassium  chloride 
the  chlorine  undergoes  ionization,  passing  over  into  anions. 
Whence  does  it  get  the  negative  charges  of  electricity  neces- 
sary to  convert  the  chlorine  atoms  into  ions?  The  only 
possible  source  is  the  platinum  electrode  immersed  in  the 
potassium  chloride.  This  electrode,  having  given  up  nega- 
tive electricity  in  converting  the  chlorine  atoms  into  chlo- 
rine ions,  is  consequently  charged  positively.  The  positive 
electricity  which  thus  accumulates  upon  this  electrode,  due 
to  the  ionization  of  the  chlorine,  flows  around  through  the 

1  Zeti.  phys.  Chem.,  9,  550  (1892). 


160  THE  NATURE  OF  SOLUTION 

galvanometer  to  the  other  electrode  immersed  in  the  solu- 
tion of  ferrous  chloride.  The  ferrous  chloride  is,  of  course, 
dissociated  by  the  water  into  ferrous  ions  each  carrying 
two  positive  charges  of  electricity,  and  into  two  chlorine 
ions  each  carrying  one  negative  charge  of  electricity.  The 
ferrous  ion  with  its  two  positive  charges  can  take  up  one 
more  positive  charge  yielding  a  ferric  ion,  which  has  three 
positive  charges  upon  it.  It  does  so,  getting  the  third 
positive  charge  from  the  platinum  electrode  immersed  in  the 
solution  of  the  ferrous  salt. 

The  ferric  ion  with  its  three  positive  charges  can  pair  off 
against  three  chlorine  ions,  each  with  one  negative  charge. 
This  is  the  condition  in  a  solution  of  ferric  chloride.  The 
question  is,  whence  comes  the  third  chlorine  ion? 

It  is  well  known  that  all  anions  move  against  the  current. 
When  the  chlorine  gas  conducted  into  the  solution  of  potas- 
sium chloride  becomes  ionized,  the  chlorine  ions  move  against 
the  current  through  the  solution  of  potassium  chloride,  and 
may  be  regarded  as  moving  through  the  siphon  over  to  the 
ferrous  chloride  and  pairing  off  against  the  ferrous  ions  which 
have  been  converted  into  ferric  ions  by  taking  positive  elec- 
tricity from  the  electrode  immersed  in  the  solution  sur- 
rounded by  the  iron  salt. 

The  fundamental  point  of  this  experiment  is,  that  the 
valence  of  a  bivalent  or  ferrous  ion  is  raised  to  a  condition 
of  trivalency,  or  to  the  ferric  condition,  by  simply  adding 
electricity  to  it,  and  doing  this  under  such  conditions  that 
we  can  say  with  reasonable  certainty  that  nothing  else  has 
taken  place.  We  have  thus  created  valence,  as  it  were,  out 
of  electrical  energy. 

From  this  experiment  alone  we  know  that  valence  is  but 
the  chemical  expression  of  the  number  of  electrons  or  free 
electrical  charges  upon  the  ions.  Thus  valence  is  placed 
upon  a  perfectly  rigid  physical  basis,  and  can  and  should  be 
discussed  primarily  in  terms  of  Faraday's  law. 

It  may  be  noted  in  passing  that  this  experiment  illus- 
trates also  another  principle,  as  was  pointed  out  by  Ostwald. 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  161 

Using  the  term  oxidation  in  the  old  sense  of  increase  hi 
valence,  the  valence  of  the  ferrous  ion  is  raised,  or  the  fer- 
rous ion  is  oxidized  by  chlorine  which  does  not  come  hi 
contact  with  it. 

This  experiment  may  then  be  regarded  as  showing  that 
contact  is  not  essential  to  chemical  action;  things  may 
react  that  do  not  touch  one  another. 

Ostwald1  has  described  another  experiment  illustrating 
chemical  action  without  mechanical  contact,  which  is  much 
more  easily  carried  out,  and  more  striking  hi  its  results  than 
the  above. 

Electrolytes  Conduct  Only  by  Undergoing  Electrolysis.  — 
It  has  long  been  known  that  when  an  electric  current  is 
passed  through  a  solution  of  an  electrolyte,  the  electrolyte 
undergoes  decomposition,  the  cations  moving  to  the  cathode 
and  the  anions  to  the  anode.  This  raises  the  question  as  to 
whether  this  is  the  only  way  hi  which  electricity  can  pass 
through  a  solution  of  an  acid,  base,  or  salt.  Might  not 
small  amounts  of  electricity  get  through  solutions  hi  some 
other  way  than  by  being  carried  by  the  ions  which  are 
simultaneously  discharged? 

This  question  was  answered  once  and  for  all  by  Ostwald 
and  Nernst2  hi  the  following  way.  A  large  balloon  flask 
was  covered  with  tin  foil,  which  was  connected  with  the 
condenser  of  a  Holtz  machine  and  charged  positively.  The 
flask  was  filled  with  dilute  sulphuric  acid,  which  was  con- 
nected by  means  of  a  siphon  with  sulphuric  acid  in  a  dish. 
Into  the  sulphuric  acid  in  the  dish  was  inverted  a  glass  tube 
drawn  out  to  a  fine  capillary  below  and  partly  filled  with 
mercury.  The  mercury  and  the  sulphuric  acid  came  hi 
contact  hi  the  capillary  tube.  The  mercury  hi  this  tube 
was  connected  with  the  earth.  When  the  frictional  ma- 
chine was  set  hi  operation,  the  tin  foil  around  the  flask 
became  charged  positively.  This  attracted  electrostatically 
the  SO4  ions  within  the  flask  coming  from  the  sulphuric 
acid,  and  repelled  the  hydrogen  ions  from  the  same  acid. 

1  Ztit.  phys.  Chem.,  9,  540  (1892).  2  Ibid.,  3,  120  (1889). 


162  THE  NATURE  OF  SOLUTION 

The  hydrogen  ions  coming  in  contact  with  the  mercury  in 
the  capillary  tube  gave  up  then*  charges  to  the  mercury  and 
separated  as  hydrogen  gas.  This  could  be  seen  in  the  fine 
capillary  by  means  of  a  microscope.  If  the  capillary  was 
carefully  calibrated,  the  amount  of  hydrogen  which  sepa- 
rated could  be  accurately  measured.  Ostwald  and  Nernst 
showed  that  when  0.000005  of  a  coulomb  of  electricity  was 
passed  through  the  solution  of  sulphuric  acid,  hydrogen  gas 
was  separated.  They  showed  that  Faraday's  law  holds  to 
within  this  limit. 

Another  experiment  devised  by  Ostwald,  bearing  on 
this  same  point,  although  only  theoretical,  should  be  re- 
ferred to  on  account  of  its  historical  interest. 

Take  two  beakers,  and  partly  fill  them  with  a  solution 
of  potassium  chloride.  Connect  the  two  with  a  siphon 
filled  with  the  solution  of  potassium  chloride,  and  bring  up 
to  one  beaker  a  condenser  charged,  we  will  say,  negatively. 
This  will  attract  electrostatically  the  positively  charged 
potassium  ions,  and  repel  the  negatively  charged  chlorine 
ions.  Now  remove  the  siphon,  when  the  beaker  next  to  the 
condenser  will  contain  an  excess  of  the  positively  charged 
potassium  ions.  Connect  the  contents  of  this  beaker  with 
the  ground,  and  the  excess  of  potassium  ions  will  lose  their 
charges  to  the  earth  and  remain  as  uncharged  potassium 
atoms.  These  would  then  react  with  water  hi  the  usual 
way  and  liberate  hydrogen  gas. 

Ostwald  stated  that  some  one  wrote  him  that  he  tried 
this  experiment,  and  failed  to  observe  the  hydrogen.  Ost- 
wald then  calculated  the  size  of  condenser  that  it  would  be 
necessary  to  use  to  liberate  enough  hydrogen  to  be  seen 
escaping  from  the  solution  of  potassium  chloride,  taking 
into  account  the  solubility  of  hydrogen  gas  in  water.  The 
condenser  if  in  the  form  of  a  cube  would  have  an  edge  a 
kilometer  long,  in  a  word,  the  condenser  would  be  a  cube 
whose  edge  would  be  about  three-fifths  of  a  mile. 

The  reason  for  all  this  will  become  obvious  when  we 
consider  the  enormous  amount  of  electricity  carried  by  a 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  163 

few  ions.  A  very  large  amount  of  electricity  is  necessary 
to  cause  enough  ions  to  separate  to  be  visible  to  the  naked 
eye.  That  is  the  reason  why  the  above  described  experiment 
of  Ostwald  and  Nernst  was  devised  to  prove  especially  to 
Du  Bois  Raymond  the  correctness  of  the  theory  of  electro- 
lytic dissociation  of  which  the  result  obtained  is  a  necessary 
consequence. 

The  Laws  of  Faraday  Apparently  Rigid  Laws  of  Nature. 
—  It  will  be  recognized  at  once  that  the  first  law  of  Faraday 
is  a  fundamental  law  of  electrolytic  conduction.  Upon  its 
rigidity  depends  our  knowledge  of  the  passage  of  electricity 
through  solutions  of  electrolytes.  Such  a  fundamental 
law  would  naturally  be  very  exhaustively  tested.  The 
strength  of  the  current  employed  to  effect  the  electrolysis 
would  be  greatly  varied.  The  temperature  would  be 
changed,  and  also  the  strength  of  the  solution  through 
which  the  current  was  passed.  All  of  these  variables  have 
been  studied  and  Faraday's  law  has  stood  the  test  as  a 
rigid  law  of  nature  would  do. 

The  rigidity  of  Faraday's  law  under  all  conditions  has 
been  called  hi  question,  and  in  the  following  way.  A  solu- 
tion electrolyzed  under  high  pressure  separated  kss  of  the 
electrolyte  for  a  given  amount  of  current,  than  when  elec- 
trolyzed under  ordinary  pressure.  This  could  be  interpreted 
as  showing  that  under  these  conditions  Faraday's  law  did 
not  hold. 

Pressure  was  brought  to  bear  on  the  solution  by  com- 
pressing the  gas  above  the  solution.  Under  these  conditions 
more  gas  would,  from  Henry's  law,  dissolve  than  when  the 
gas  was  under  only  normal  pressure.  It  was  shown  that  a 
small  part  of  this  dissolved  gas  was  ionized  and  this  assisted 
hi  carrying  the  current  through  the  solution.  When  this 
was  taken  into  account  it  was  found  that  the  law  of  Faraday 
held  also  under  this  condition. 

The  Laws  of  Faraday  are  to  be  placed  among  the  few 
rigid  laws  of  nature  to  which  no  exception  is  thus  far  known, 
and  the  number  of  such  laws  is  very  few  indeed. 


164  THE  NATURE  OF  SOLUTION 

The  Nature  of  Electrolysis.  —  When  a  continuous  current 
is  passed  through  a  dilute  solution  of  an  acid,  base,  or  salt, 
hydrogen  is  liberated  at  the  cathode  and  oxygen  at  the 
anode.  If  we  are  dealing  with  the  salt  of  a  metal  which  does 
not  act  on  water  the  metal  itself  is  deposited  at  the  cathode. 
As  we  have  seen,  the  only  way  an  electric  current  can  pass 
through  a  solution  or  through  water  is  by  electrolyzing  the 
dissolved  substance  or  the  solvent.  Such  are  some  of  the 
facts  of  electrolysis.  What  is  then'  explanation? 

We  have  become  familiar  with  the  theories  proposed  by 
Grotthuss  and  by  Clausius  to  account  for  the  phenomena  of 
electrolysis.  The  theory  of  the  electrolysis  of  acids,  bases, 
and  salts,  which  was  held  for  forty  years,  was  based  upon 
the  suggestions  made  by  Clausius.  The  theory,  as  it  was 
applied,  is  as  follows. 

Older  Theory.  —  When  the  current  is  passed  through  a 
solution  of  an  acid,  the  hydrogen  ions  of  the  acid  move  over 
to  the  cathode,  give  up  their  positive  charges,  and  escape  as 
hydrogen  gas.  The  anion  moves  over  to  the  anode,  gives 
up  its  charge,  but  not  being  volatile  does  not  escape.  It 
"acts  on  water,"  as  it  was  said,  liberating  oxygen  and  com- 
bining with  the  hydrogen,  reforming  the  acid. 

The  electrolysis  of  a  base  was  strictly  analogous.  The 
hydroxyl  anions  move  over  to  the  anode,  give  up  their 
charges,  two  of  them  react  with  one  another,  forming  a 
molecule  of  water  and  liberating  oxygen  which  escapes  at 
the  anode.  The  cation  of  the  base  moves  over  to  the 
cathode,  gives  up  its  charge,  and  then  reacts  with  water, 
reforming  the  base  and  liberating  hydrogen  which  escapes. 

In  the  electrolysis  of  a  salt,  the  cation  moves  over  to  the 
cathode,  gives  up  its  positive  charge,  then  reacts  with 
water  liberating  hydrogen  gas  which  escapes.  The  anion 
moves  over  to  the  anode,  gives  up  its  charge,  and  reacts 
with  water  liberating  oxygen. 

It  is  a  very  simple  matter  to  show  that  this  older  theory 
is  untenable.  Take  the  case  of  an  acid.  The  anion  goes 
over  to  the  anode,  gives  up  its  charge,  and  then  "reacts 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  165 

with  water."  Just  what_does  this  mean?  In  the  case 
of  nitric  acid,  the_anion,  NO3,  finds  around  the  anode  hy- 
droxyl  anions  (OH)  from  the  dissociated  water.  Which 
will  give  up  its  charge?  Obviously  the  one  that  holds  it 
less  firmly.  The  old  theory  said  that  the  NO3  gives  up 
its  charge  because  it  holds  it  less  firmly  than  the  hydroxyl, 
than  takes  the  charge  from  the  Jiydroxyl,  which  must 
hold  its  charge  more  firmly  than  NO3,  otherwise  hydroxyl 
and  not  NO3  would  have  given  up  its  charge.  This  is,  of 
course,  a  reductio  ad  absurdum. 

The, same  applies  to  the  older  theory  of  the  electrolysis 
of  a  base.  The  cation  going  to  the  cathode  finds  there 
hydrogen  cations  from  the  dissociated  water.  The  cation 
of  the  base  gives  up  its  charge  and  not  the  hydrogen,  because 
the  latter  holds  its  charge  more  firmly  than  the  former.  The 
cation  of  the  base,  having  been  discharged,  takes  the  charge 
from  the  hydrogen,  which  must  have  held  its  charge  more 
firmly  than  the  cation  of  the  base,  otherwise,  the  hydrogen 
and  not  the  cation  of  the  base  would  have  given  it  up 
originally.  The  same  absurdity  manifests  itself  as  hi  the 
case  of  an  acid. 

With  a  salt  the  old  theory  leads  to  an  absurdity  at  both 
the  cathode  and  the  anode.  The  cation  of  the  salt  goes 
to  the  cathode  and  gives  up  its  charge  because  it  holds  the 
charge  less  firmly  than  the  hydrogen  ions  from  the  disso- 
ciated water.  It  then  reacts  with  water,  which  means  that 
it  takes  the  charge  from  the  hydrogen  ions,  which  is,  of 
course,  impossible.  The  anion  of  the  salt  goes  to  the  an- 
ode, gives  up  its  charge,  because  it  holds  it  less  firmly  than 
the  hydroxyl  anion  from  the  dissociated  water,  then  takes 
the  charge  from  the  hydroxyl  anion  —  which  is,  again,  im- 
possible. 

It  is  thus  comparatively  simple  to  show  that  the  older 
theory  of  electrolysis,  based  upon  the  views  of  Clausius,  is 
hi  error.  It  is  a  very  different  matter  to  show  what  the 
correct  view  is.  This,  however,  has  been  done. 

The  Decomposition  Values  of  the  Ions.  —  It  will  be  seen 


166  THE  NATURE  OF  SOLUTION 

from  the  above  that  the  fundamental  question  is,  which 
ions  hold  their  charges  more  firmly,  the  hydrogen  and  hy- 
droxyl  from  the  dissociated  water,  or  the  anions  and  cations 
of  acids,  bases,  and  salts?  The  question  has  been  answered 
directly  and  very  satisfactorily  by  Le  Blanc,1  who  meas- 
ured the  so-called  discharging  values  of  the  ions.  What  does 
this  mean?  The  discharging  value  of  an  ion  is  the  voltage 
which  is  necessary  and  just  sufficient  to  cause  the  ions  in 
question  to  lose  then-  charge.  This  is  obviously  the  voltage 
which  is  necessary  and  just  sufficient  to  cause  a  continuous 
current  to  flow  through  the  solution  containing  the  ions  in 
question. 

The  results  that  were  obtained  will  be  given  hi  connection 
with,  or  as  necessary  consequences  of,  the  theory  which  was 
proposed  to  account  for  them. 

Newer  Theory  of  Electrolysis.  —  In  the  electrolysis  of 
an  add,  the  hydrogen  cations,  together  with  the  hydrogen 
ions  from  the  dissociated  water,  move  to  the  cathode, 
give  up  their  charges  and  escape.  The  anions  of  the  acid 
move  to  the  anode  and  find  around  this  pole  hydroxyl 
anions  from  the  water.  Which  will  discharge?  Obviously 
the  one  which  holds  its  charge  less  firmly.  Le  Blanc  has 
shown  that  hydroxyl  has  a  lower  discharging  value  than 
almost  any  other  anion  except  perhaps  chlorine,  bromine, 
and  the  like.  Therefore,  whenever  we  electrolyze  almost 
any  acid  other  than  the  halogen  acids,  only  oxygen  is  given 
off  at  the  anode.  The  hydroxyls  from  the  dissociated  water 
discharge  at  the  anode,  two  of  them  combine  and  liberate 
oxygen  in  the  usual  way,  and  the  anion  of  the  acid  simply 
remains  paired  off  against  a  hydrogen  ion  which  may  have 
come  either  from  the  dissociated  water  or  from  the  acid. 

The  consequences  of  this  are  very  interesting  and  can 
readily  be  tested  by  the  results  of  experiment. 

Newer  Theory  and  Decomposition  Values.  —  If  the 
above  suggestion  is  correct,  the  decomposition  values  of 
solutions  of  the  strong  acids  having  the  same  concentrations 

*  Zeit.  phys.  Chem.,  8,  299  (1891). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  167 

must  be  constant.  The  reason  for  this  is  obvious.  The 
decomposition  values  of  such  solutions  of  such  acids  are 
simply  the  voltages  which  are  just  necessary  to  decompose 
hydrogen  and  hydroxyl  ions  of  the  concentrations  at  which 
they  exist  hi  the  solutions  hi  question.  But  since  these 
solutions,  being  of  the  same  concentrations,  have  the  same 
concentrations  of  both  hydrogen  and  hydroxyl  ions,  we  have 
in  them  the  same  numbers  of  the  same  lands  of  ions,  and  they 
must,  of  course,  have  the  same  decomposition  values.  The 
following  results  show  that  the  above  conclusion  is  in  accord 
with  the  facts.  These  results  were  obtained  for  normal 
solutions.1 

Decomposition 

Acid  values 

d-Tartaric  1.62  volts 

Perchloric  1.65     " 

Malonic  1.69     " 

Dichloracetic  1.66 

Monochloracetic  1.72     " 

Nitric  1.69     "! 

Sulphuric  1.67     "' 

The  approximate  constancy  of  these  decomposition  values 
is  in  keeping  with  the  above  conclusion. 

In  the  case  of  bases,  the  new  theory  says  that  the  hy- 
droxyl anions  move  to  the  anode,  give  up  their  charge  and 
react  with  one  another  forming  water  and  liberating  oxygen. 
Only  hydroxyl  ions  discharge  at  the  anode.  Most  of  these 
come  from  the  base,  a  few  from  the  dissociated  water. 

The  cation  of  the  base  goes  to  the  cathode  and  finds  there 
hydrogen  ions  from  the  dissociated  water.  Le  Blanc  has 
shown  that  hydrogen  has  a  lower  discharging  value  than 
most  of  the  cations  of  bases.  Therefore  hydrogen  is  lib- 
erated. The  salts  of  certain  metals,  such  as  copper,  yield 
a  cation  with  a  low  discharging  value,  and  hi  such  cases,  hi 
addition  to  the  hydrogen  liberated,  we  have  the  metal 
separating  on  the  cathode. 

What  has  been  said  above,  however,  holds  hi  general. 
The  hydrogen  ions  are  discharged  at  the  cathode.  If  we 

1  Zett.  phys.  Chem.,  8,  315  (1891). 


168  THE  NATURE  OF  SOLUTION 

use  the  same  concentrations  of  the  different  strong  bases, 
we  have  approximately  the  same  concentrations  of  hydrogen 
ions  around  the  cathode,  and  hydroxyl  ions  around  the 
anode,  regardless  of  the  base  that  we  use.  The  electrolysis 
of  normal  solutions  of  bases  is  therefore  the  decomposition 
of  hydrogen  and  hydroxyl  ions  at  the  concentrations  at 
which  they  exist  in  these  solutions.  The  decomposition 
values  of  normal  solutions  of  bases  should  therefore  be  a  con- 
stant. This  conclusion  is  also  readily  tested  by  the  results 
obtained  by  Le  Blanc.1 

Decomposition 

Base  value 

Sodium  hydroxide  1.69  volts 

Potassium  hydroxide  1.67 

Ammonium  hydroxide  1.74     " 

|  Methylamine  1.75     " 

|  Diethylamine  1.68     " 

5  Tetraethyl  ammonium  hydrate  1.74     " 

o  . 

The  conclusion  is  borne  out  by  the  facts. 

One  other  relation  should  be  pointed  out  while  con- 
sidering acids  and  bases.  In  a  normal  solution  of  a  strong 
acid  the  concentration  of  the  hydrogen  ions  is  the  same  as 
the  concentration  of  the  hydroxyl  ions  in  a  normal  solution 
of  a  strong  base.  The  hydrogen  ions  in  the  acid  are  equal 
in  number  to  the  hydroxyl  ions  in  the  base. 

The  hydroxyl  ions  in  the  acid  solution  coming  from  the 
dissociated  water  are  equal  in  number  to  the  hydrogen 
ions  in  the  solution  of  the  base  which  also  come  from  the 
dissociated  water.  Since  both  solutions  contain  the  same 
number  of  hydrogen  ions  and  the  same  number  of  hydroxyl 
ions,  the  discharging  values  of  the  two  must  be  the  same,  and 
the  above  results  show  this  to  be  the  case. 

Let  us  now  consider  the  electrolysis  of  salts.  The 
cation  of  the  salt  moves  to  the  cathode,  and  the  hydrogen 
there  from  the  dissociated  water  discharges.  The  cation 
remains  around  the  cathode,  paired  off  against  the  hydroxyl 

1  ZeU.  phys.  Chem.,  8,  315  (1891). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  169 

ion  which  comes  from  the  same  molecule  of  water  as  the 
hydrogen  ion.  The  anion  of  the  salt  moves  to  the  anode, 
the  hydroxyl  there  from  the  dissociated  water  discharging. 
The  anion  remains  paired  off  against  the  hydrogen  ion  which 
came  from  the  same  molecule  of  water  as  the  hydroxyl  ion 
which  has  discharged.  There  is  thus  an  accumulation  of 
hydroxyl  anions  around  the  cathode  and  of  hydrogen  cations 
around  the  anode.  What  will  be  the  effect  of  this  on  the 
dissociation  of  water? 

There  is  a  well-known  general  principle,  that  the  presence 
of  any  ion  drives  back  the  dissociation  of  any  electrolyte 
yielding  that  ion.  Therefore,  the  presence  of  hydroxyl  ions 
around  the  cathode  and  of  hydrogen  ions  around  the  anode, 
drives  back  or  diminishes  the  dissociation  of  the  water  around 
both  of  these  poles.  There  being  a  smaller  number  of  ions 
to  discharge,  it  will  require  a  higher  voltage  to  discharge 
them,  the  decomposition  values  depending  not  only  on  the 
nature  of  the  ions,  but  also  on  their  number. 

Salts  should  therefore  have  higher  discharging  or  decom- 
position values  than  either  acids  or  bases.  We  see  that  such 
is  the  case.1 

Decomposition 

Salts  values 

Potassium  sulphate  2.20  volts 

Sodium  sulphate  2.21 

Potassium  nitrate  2.17 

Sodium  nitrate  2.15 

Potassium  chloride  1.96 

Sodium  chloride  1.98 

Ammonium  nitrate  2.08 

Ammonium  sulphate  2.11 

Calcium  chloride  1.89 

Calcium  nitrate  2.11 

The  facts  are  hi  keeping  with  the  conclusions  from  theory. 

Thus  we  see  that  all  of  the  conclusions  from  the  above 
theory  of  electrolysis  are  borne  out  by  the  experimental 
results.  We  can,  therefore,  regard  this  theory  as  experi- 
mentally verified,  and  can  accept  it  tentatively  until  some 
one  suggests  another  theory  which  harmonizes  better  with 
the  facts. 

1  Le  Blanc:  Zeit.  phys.  Chem.,  8,  311  (1891). 


170  THE  NATURE  OF  SOLUTION 

Electrolysis  of  Water  a  Direct  Decomposition  by  the 
Current.  —  It  may  at  first  sight  seem  that  there  is  really 
very  little  difference  between  the  older  theory  of  electrolysis 
which  is  now  regarded  as  untenable,  and  the  newer  view 
supported  by  the  decomposition  values  found  by  Le  Blanc. 
A  moment's  thought  will  show  that  such  is  not  the  case. 

The  older  theory  maintained  that  the  decomposition  of 
water  in  electrolysis  was  not  the  direct  action  of  the  current, 
but  was  a  secondary  act.  The  direct  products  of  electrolysis, 
the  discharged  cations  and  anions  of  electrolytes,  reacted 
with  the  water  decomposing  it,  liberating  hydrogen  at  one 
pole  and  oxygen  at  the  other.  The  decomposition  of  the 
water,  in  terms  of  the  older  theory  of  electrolysis,  was 
then  a  secondary  act,  not  brought  about  by  the  current 
at  all,  but  by  the  products  of  the  direct  action  of  the 
current. 

The  newer  theory  of  electrolysis  is  exactly  the  reverse. 
The  hydrogen  ions  of  the  dissociated  water  are  discharged 
at  the  cathode,  and  the  hydroxyl  ions  of  the  dissociated 
water  at  the  anode.  The  water  is  decomposed  directly  by 
the  current,  and  this  theory  of  electrolysis  is  sometimes  known 
as  the  theory  of  the  primary  decomposition  of  water. 

The  discharging  values  of  the  ions  are  also  the  basis  of 
electrolytic  separation  of  the  metals,  which  is  a  new  and 
comprehensive  branch  of  analytical  chemistry,  but  it  would 
lead  us  too  far  to  discuss  it  hi  any  detail  here. 

PROPERTY  OF  SOLUTIONS  OF  ELECTROLYTES  TO  CONDUCT 
THE  ELECTRIC  CURRENT 

Solutions  of  electrolytes  differ  from  those  of  non-elec- 
trolytes in  that  the  former  conduct  the  current  while  the 
latter  do  not.  Indeed,  this  is  the  fundamental  distinction 
between  these  two  great  classes  into  which  all  chemical 
compounds  fall. 

The  power  of  acids,  bases,  and  salts  to  conduct  the  cur- 
rent in  aqueous  solution  must  be  regarded  as  among  the 
most  fundamental  and  important  properties  of  solutions  of 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  171 

these  substances.  It  not  only  distinguishes  the  electrolytes 
from  the  non-electrolytes,  but  the  study  of  this  property  has 
done  much  to  throw  light  on  the  nature  of  solution  in  gen- 
eral. Some  stress  will  therefore  be  laid  on  the  measurements 
of  conductivity,  and  on  the  discussion  of  the  results  obtained. 

Principle  Involved  in  Measuring  Conductivity.  —  In 
determining  the  conductivity  of  a  piece  of  metal,  we  must 
define  the  size  and  shape  of  the  piece  which  we  would  study, 
since  resistance,  and  consequently  conductivity,  is  a  function 
of  the  diameter  and  length,  as  well  as  of  the  nature  of  the 
substance  itself.  In  the  case  of  metals,  we  may  take  a  cube 
of  the  metal  whose  edge  is  a  centimeter  hi  length;  but 
since  this  would  have  very  small  resistance,  and,  therefore, 
great  conductivity,  it  is  better  to  take  a  cylinder  of  the 
metal  one  meter  hi  length  and  with  a  cross-section  of  one 
square  milhmeter.  The  latter  would  have  100xlOO  =  104 
times  the  resistance  of  the  former,  and  therefore,  only  one 
ten  thousandth  of  the  conductivity  of  the  cubic  centimeter 
of  metal. 

In  the  case  of  solutions  we  must  choose  some  standard  for 
the  solution  to  be  measured.  The  conductivities  of  solu- 
tions are  far  less  than  those  of  metals,  so  we  take  as  a 
standard  the  conductivity  of  a  cube  of  the  solution  whose 
edge  is  one  centimeter. 

In  order  that  the  results  for  different  substances  should 
be  comparable,  we  must  deal  with  comparable  numbers  of 
molecules  or  the  ions  resulting  therefrom,  and  we  must  be 
able  to  measure  the  conductivities  of  solutions  of  any  and 
all  concentrations. 

The  conductivity  of  the  cube  of  solution  referred  to  above 
we  will  call  the  specific  conductivity  and  will  represent  this 
by  c.  If  we  represent  the  number  of  cubic  centimeters  of  the 
solution  in  question  which  contain  a  gram-molecular  weight 
of  the  substance  by  p,  then  the  molecular  conductivity,  which 
is  sometimes  called  X,  but  more  frequently  /*,  will  be 

X  or  p,  —  pc 


172  THE  NATURE  OF  SOLUTION 

The  Kohlrausch  Method  of  Measuring  Electrical  Con- 
ductivity. —  To  determine  the  molecular  conductivity,  ju,  of 
any  solution,  we  must  be  able  to  determine  c,  or  its  con- 
ductivity, which  is  the  reciprocal  of  its  resistance,  the 
quantity  measured.  Kohlrausch  has  devised  a  very  satis- 
factory method  for  measuring  the  conductivity  of  solutions. 

When  a  continuous  current  is  passed  through  a  solution, 
electrolysis  results,  hydrogen  separating  at  one  pole  and 
oxygen  at  the  other.  The  poles  become  covered  with  gases, 
or  they  become  polarized,  as  we  say.  Kohlrausch  overcame 
this  difficulty  by  using  not  a  continuous,  but  an  alternating 
current.  Such  a  current  from  a  small  induction  coil  was 
sent  through  the  solution  placed  in  a  glass  cup  containing 
two  platinum  electrodes,  thrown  into  one  arm  of  a  wheat- 
stone  bridge,  and  a  rheostat  or  resistance  box  thrown  into 
the  other  arm.  The  balance  was  established  by  means  of  a 
telephone  receiver  placed  one  arm  between  the  solution  and 
the  rheostat,  and  the  other  attached  to  the  bridge  by  a 
sliding  contact  which  is  moved  along  the  bridge  wire  until 
the  hum  of  the  inductorium  is  no  longer  heard  in  the  tele- 
phone. Instead  of  a  telephone  receiver  a  dynamometer 
or  alternating  current  galvanometer  may  be  employed. 

The  calculation  of  the  resistance  offered  by  the  solution 
is  very  simple  indeed.  If  we  represent  the  conductivity 
referred  to  molecular  quantities  as  is  usually  done  by  jit,,  v 
being  the  volume  of  the  solution  or  the  number  of  liters  of 
the  solution  that  contain  a  gram-molecular  weight  of  the 
acid,  base,  or  salt;  we  have  from  the  principle  of  the  wheat- 
stone  bridge 


K  being  the  constant  of  the  cell  which  depends  for  its  value 
on  the  size  of  the  electrode  plates  and  the  distance  they  are 
apart,  m  and  n  the  readings  on  the  bridge  and  A  the  resis- 
tance in  ohms  in  the  rheostat.1 

1  For  further  details  in  connection  with  the  deduction  of  this  equation 
see  the  Author's  Elements  of  Physical  Chemistry,  4th  edition,  p.  392  (The 
Macmillan  Co.). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  173 

Applying  the  Kohlrausch  Conductivity  Method.  —  There 
are  two  quantities  in  the  conductivity  equation  which 
cannot  be  determined  directly  by  separate  experiments. 
These  are  ju,  or  the  molecular  conductivity  which  is  the 
value  desired,  and  K  the  cell  constant.  Whenever  we  have 
only  one  equation  and  two  unknown  quantities  we  must 
by  some  means  eliminate  one  of  the  unknowns.  In  this 
case  we  eliminate  K,  i.e.,  the  cell  constant. 

This  was  done  for  the  first  cell  by  measuring  the  size 
of  the  plates  and  their  distance  apart.  Now  that  we  know 
the  value  of  ju  for  many  substances  it  is  not  necessary  to 
proceed  in  this  way.  To  " standardize  "  a  cell  or  determine 
the  "cell  constant/'  we  now  proceed  as  follows.  We  take  a 
solution  of  some  substance  whose  molecular  conductivity 
at  the  dilution  and  temperature  is  known.  We  place  it 
hi  the  cell,  bring  it  to  the  desired  temperature,  and  then 
determine  the  other  values  hi  the  equation,  the  lengths  of 
the  bridge  arms,  and  the  resistance  in  the  rheostat,  by  direct 
readings.  We  now  have  only  one  unknown  hi  the  equation 
—  the  cell  constant  K,  and  we  solve  for  this.  We  usually  use 

a  —  solution  of  potassium  chloride,  which  has  a  molecular 

oU 

conductivity  of  129.7  Siemen's  units  or  137.9  ohms  (recip- 
rocal ohms)  at  25°. 

The  question  of  the  purity  of  the  water  used  hi  preparing 
the  solutions  whose  conductivities  are  to  be  measured  is 
fundamental.  What  we  actually  measure  is  the  conductivity 
of  the  solution  plus  that  of  the  water,  and  we  must  make 
the  latter  as  small  as  possible.  We  then  subtract  the  con- 
ductivity of  the  water  from  the  total  conductivity  found, 
and  the  remainder  is  the  conductivity  of  the  solution  hi 
question. 

It  might  seem  from  this  that  the  actual  value  of  the  con- 
ductivity of  the  water  used  was  of  very  little  significance. 
It  would  be  only  necessary  to  determine  the  conductivity 
of  the  water,  whatever  it  might  be,  and  then  subtract  this 
from  the  total  conductivity,  in  order  to  get  the  conductivity 


174  THE  NATURE  OF  SOLUTION 

of  the  solution.  A  moment's  thought  will  show  that  this 
will  not  do.  Suppose  we  are  measuring  the  conductivity  of 
an  acid,  and  the  water  used  contains  ammonia.  We  are 
then  really  measuring  to  some  degree,  not  the  conductivity 
of  the  acid,  but  of  its  ammonium  salt.  The  conductivity 
of  the  ammonium  salt  of  any  acid  is  only  a  small  part  of 
the  conductivity  of  the  acid  itself,  because  the  ammonium 
ion  moves  far  more  slowly  than  the  hydrogen  ion.  The 
values  that  we  would  thus  obtain  would  bear  no  direct  re- 
lation to  the  values  that  we  wish  to  secure. 

If  the  substance  whose  conductivity  we  were  measuring 
was  a  base,  and  the  water  used  contained  carbon  dioxide,  we 
would  to  some  extent  be  measuring  the  conductivity  of  the 
carbonate  of  the  base.  The  carbonate  would_have  far  less 
conductivity  than  the  free  base,  because  the  COs  ion  moves 
so  much  more  slowly  than  the  hydroxyl  ion.  The  results 
obtained  would,  therefore,  have  no  real  significance. 

Purification  of  Water.  —  Pure  water  has  never  been 
prepared.  For  that  matter,  nothing  has  ever  been  prepared 
pure,  and  nothing  ever  can  be  while  the  present  laws  of 
nature  obtain;  since  everything  is  soluble  in  everything  else 
with  which  it  comes  in  contact.  The  purest  water  has 
been  prepared  by  Kohlrausch  and  Heydweiller,1  and  in  the 
following  way.  Water  which  had  been  purified  by  distilla- 
tion was  subjected  to  fractional  crystallization,  and  the 
purification  carried  as  far  as  possible  by  this  process.  It 
was  then  placed  in  one  arm  of  a  {/-shaped  platinum  vessel 
which  was  attached  to  an  air-pump.  After  removing  the 
ah*,  a  part  of  the  water  was  distilled  from  one  arm  of  the 
vessel  over  into  the  other.  Into  the  receiving  arm  electrodes 
had  been  inserted,  and  the  conductivity  of  the  water  deter- 
mined without  allowing  it  to  come  in  contact  with  the  air. 

The  easiest  way  to  remember  the  value  of  the  resistance 
or  conductivity  of  water  thus  purified,  is  as  follows.  The 
electrical  resistance  of  a  millimeter  cube  of  this  water  at 
zero  is  the  same  as  that  of  a  copper  wire  whose  cross- 

1  Zett.  phys.  Chem.,  14,  317  (1894). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  175 

section  is  a  square  millimeter,  the  length  of  the  wire  being 
such  that  it  could  be  wrapped  around  the  earth  at  the  equator 
one  thousand  times.  Such  water  would  be  ideal  for  con- 
ductivity work,  but  it  is  next  to  impossible  to  obtain  water 
of  this  degree  of  purity  in  any  quantity,  and  further,  if  such 
water  could  be  obtained  hi  quantity,  it  could  not  be  preserved 
with  this  degree  of  purity  during  use  for  conductivity  pur- 
poses. 

We  must  then  devise  a  method  for  obtaining  water  of 
purity  sufficient  for  conductivity  work,  and  hi  fairly  large 
quantity.  A  number  of  such  methods  have  been  devised. 
The  one  used  in  this  laboratory  consists  hi  distilling  ordi- 
nary distilled  water  from  chromic  acid,  and  then  redis- 
tilling the  product  from  barium  hydroxide.  The  chromic 
acid  burns  up  the  organic  matter  hi  the  water  and  holds 
back  the  ammonia.  The  carbon  dioxide  which  passes  over 
is  "fixed"  by  the  barium  hydroxide,  and  the  product  is 
pure  enough  for  conductivity  purposes.  By  this  method 
about  ten  liters  of  water  a  day  may  be  obtained,  having 
a  specific  conductivity  of  about  1  X  10  ~6. 

Regulation  of  Temperature  —  We  shall  see  that  the  con- 
ductivities of  solutions  of  electrolytes  increase  enormously  with 
rise  in  temperature.  This  is  exactly  the  reverse  of  what  takes 
place  with  metals.  The  conductivities  of  the  latter  decrease 
with  rise  in  temperature,  or  increase  with  fall  in  temperature; 
becoming  infinite  at  or  near  .absolute  zero,  as  has  recently 
been  shown  experimentally  by  Kammerlingh  Onnes. 

The  temperature  coefficients  of  conductivity  of  solutions 
of  electrolytes  are  as  much  as  several  percent  per  degree  rise 
in  temperature.  Therefore  the  temperature  must  be  very 
accurately  regulated  hi  making  conductivity  measurements. 

A  large  number  of  regulating  devices  have  been  proposed 
for  this  purpose.  The  Ostwald l  regulator  did  good  service 
hi  its  day.  Better  forms,  however,  are  now  available.2 

1  Zeit.  phys.  Chem.,  2,  565  (1888). 

2  Carnegie  Institution  of  Washington,   Publication   No.    198,  Chap.   Ill 
(1914);  No.  210,  Chap.  VI  (1915).    Journ.  Amer.  Chem.  Soc. 38,  516  (1916). 


176  THE  NATURE  OF  SOLUTION 

Some  Results  of  Conductivity  Measurements.  —  An 
examination  of  the  conductivity  results  thus  far  obtained 
will  show  that,  with  respect  to  their  power  to  carry  the 
electric  current,  electrolytes  fall  into  three  classes.  First, 
strong  acids,  which  have  the  greatest  conductivities;  second, 
strong  bases,  which  have  intermediate  conductivities;  and 
third,  salts,  which  have  the  smallest  conductivities.  We 
can  readily  understand  why  the  electrolytes  fall  in  this  order. 
The  acids  yield  the  hydrogen  ion,  which  is  the  swiftest  of  all 
the  ions  aijd  therefore  gives  the  largest  conductivity.  The 
bases  yield  the  hydroxyl  ion,  and  this,  next  to  the  hydrogen, 
is  the  swiftest  of  all  the  ions.  The  ions  of  the  salts  move 
with  very  much  smaller  velocities  than  the  hydrogen  and 
the  hydroxyl  ions,  and,  therefore,  the  conductivities  of  the 
salts  are  very  much  less  than  those  of  the  strong  acids  and 
the  strong  bases. 

Among  the  acids  very  different  degrees  of  conductivity 
are  represented.  The  strongest  acids  are  the  best  conductors 
of  all  the  electrolytes,  as  has  been  stated;  but  we  have  all 
degrees  of  conductivity  represented  by  the  different  acids. 
We  have  aqueous  solutions  of  such  acids  as  hydrocyanic, 
carbonic,  boric,  which  are  scarcely  electrolytes  at  all.  We 
have  solutions  of  the  organic  acids  which  have  conductivi- 
ties all  the  way  from  very  small  to  very  large  values;  the 
strongest  organic  acids  having  conductivities  of  the  same 
order  of  magnitude  as  the  strong  mineral  acids. 

Among  the  bases  we  also  have  many  degrees  of  conduc- 
tivity represented.  We  have  many  examples  of  very  weak 
bases  such  as  ammonia  and  its  derivatives.  We  also  have 
the  very  strong  bases  —  the  hydroxides  of  the  alkalies,  and 
also  bases  with  intermediate  conductivities.  None  of  the 
salts  are  as  good  conductors  as  the  strongest  acids  and  bases, 
but  they  are,  in  general,  good  conductors  of  electricity. 
Nearly  all  of  the  salts  are  strongly  dissociated  compounds, 
and  therefore  yield  in  aqueous  solution  a  large  number  of 
ions.  Even  if  these  ions  do  not  move  with  such  velocities  as 
the  hydrogen  and  hydroxyl  ions,  their  large  numbers  give 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  177 

the  solutions  high  conductivities.  There  are  a  few  excep- 
tions to  the  aqueous  solutions  of  the  salts  being  good  con- 
ductors. The  halogen  salts  of  mercury  hi  aqueous  solution 
are  nearly  non-electrolytes.  These  solutions  have  almost 
no  conductivity  at  all.  The  same  applies  hi  a  much  less 
degree  to  the  halides  of  cadmium  and  zinc.  The  meaning 
of  these  facts  is  at  present  entirely  unknown. 

For  any  given  electrolyte  the  molecular  conductivity 
usually  increases  with  the  dilution  of  the  solution  until  a 
certain  maximum  value  is  reached.  From  this  point  on  the 
molecular  conductivity  has  the  maximum,  constant  value. 

There  are,  however,  many  exceptions  to  this  statement, 
as  Sachkanov1  has  shown.  Certain  substances,  hi  solvents 
with  small  dielectric  constants,  show  an  increase  hi  the  molec- 
ular conductivity  with  increase  hi  the  concentration  of  the 
solution.  This  is  accounted  for  by  Sachanov  as  due  to  the 
polymerization  of  the  solute  in  such  solvents,  especially  in 
the  more  concentrated  solutions,  with  the  formation  of  com- 
plex molecules  which  are  more  highly  dissociated  than  the 
simpler  molecules. 

Law  of  Electrical  Conductivity.  —  The  fundamental  law 
of  the  electrical  conductivity  of  aqueous  solutions  of  elec- 
trolytes was  discovered  by  Kohlrausch,  who  devised  the 
method  now  almost  universally  used  for  such  work.  The 
general  statement  of  the  law  is  this.  Each  ion  moves  with 
its  own  definite  velocity  which,  in  any  given  solvent,  and  under 
a  constant  driving  force,  is  constant.  This  velocity  does  not 
depend  on  the  nature,  number,  or  condition  of  the  other 
ions  present  in  the  same  solution.  If  we  represent  the  maxi- 
mum molecular  conductivity  for  any  substance,  already 
referred  to,  by  /j^,,  the  molecular  conductivity  at  infinite 
dilution,  we  would  have  — 


Kc  being  a  constant  dependent  for   its   value  upon  the 
nature  of  the  cation,  and  Ka  a  constant  dependent  for  its 

1  Zeit.  Elektrochem,  20,  529  (1914). 


178  THE  NATURE  OF  SOLUTION 

value  upon  the  nature  of  the  anion.  Ostwald  has  gener- 
alized this  law  of  Kohlrausch,  extending  it  from  solutions  of 
infinite  dilution  to  solutions  of  all  dilutions.  The  difference 
between  a  very  dilute  solution  and  one  more  concentrated, 
from  the  present  standpoint,  is  that  the  former  is  completely 
dissociated  while  the  latter  is  only  partly  broken  down  into 
ions.  In  attempting  to  apply  this  law  to  solutions  in  gen- 
eral, the  dissociations  of  these  solutions  must  be  taken  into 
account;  and  this  is  just  what  Ostwald  did.  He  showed 
that  if  the  dissociation  of  the  solution  in  question  is  multi- 
plied into  the  above  equation,  the  Kohlrauseh  law  can  be 
applied  to  solutions  of  electrolytes  in  general.  Calling  the 
percentage  of  dissociation  a,  we  have, 


Dissociation  of  Electrolytes  Measured  by  the  Conductivity 
Method.  —  The  chief  scientific  use  of  the  conductivity 
method  is  to  measure  electrolytic  dissociation.  This  appli- 
cation is  theoretically  very  simple.  If  the  solution  is  not 
dissociated  it  does  not  conduct.  If  the  dissociation  is  com- 
plete, the  conductivity  is  greater  than  at  any  other  dilution, 
as  we  are  dealing  with  molecular  quantities.  If  the  dissocia- 
tion is  somewhere  between  zero  and  one  hundred  percent, 
the  conductivity  of  the  solution  will  be  between  zero  and 
its  greatest  value.  Since  conductivity  is  proportional  to 
dissociation,  the  dissociation  at  any  dilution  is  simply  the 
ratio  between  the  molecular  conductivity  at  the  dilution 
hi  question  and  the  maximum  value  of  the  molecular  con- 
ductivity. If  we  represent  this,  as  is  usually  done  by  JJL^, 
we  have  the  dissociation  a.  : 

m£& 

/loo 

To  find  IJL^  for  strong  electrolytes,  it  is  only  necessary  to 
increase  the  dilution  of  the  solution  until  complete  dis- 
sociation is  reached,  and  then  determine  the  molecular 
conductivity  of  such  a  solution  which  is  the  value  of  /*„ 
desired. 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  179 

In  case  the  compound  in  question  is  only  weakly  dis- 
sociated, as  with  weak  acids  and  weak  bases,  we  cannot 
determine  /z^  directly  by  the  method  just  discussed.  The 
dilution  of  the  solution  hi  which  nm  is  reached  is  so  great, 
that  the  conductivity  method  cannot  be  applied  to  it.  In 
such  a  solution  the  solvent  would  conduct  more  than  the 
dissolved  substance. 

For  such  substances  an  indirect  method1  of  determining 
fj,m  has  been  worked  out,  which  is  very  efficient,  but  which 
would  lead  us  too  far  to  discuss  it  in  any  detail. 

Relation  Between  Dissociation  and  Dilution.  —  It  has 
been  known  from  the  tune  that  the  dissociation  theory 
was  proposed,  that  the  dissociation  of  solutions  increases 
with  the  dilution  of  those  solutions.  There  is  a  dilution  at 
which  any  electrolyte  is  completely  dissociated.  This 
knowledge  was  purely  qualitative.  What  we  want  to  know 
is,  how  much  does  the  dissociation  of  a  solution  increase 
for  a  given  increase  hi  its  dilution?  What  is  the  relation 
between  the  rate  of  increase  hi  the  dissociation  and  the  rate 
of  increase  hi  the  dilution? 

This  question  was  first  answered  by  Ostwald,  hi  what  has 
come  to  be  known  as  his  dilution  law.  Connecting  these 
two  quantities,  dissociation  and  dilution,  he  deduced  the 
following  expression.2 

a2 


(1  -  a)v 


constant 


hi  which  a.  is  the  percentage  dissociation  of  the  solution, 
and  v  the  volume  of  the  solution,  or  number  of  liters,  that 
contains  a  gram-molecular  weight  of  the  electrolyte. 

This  dilution  law  of  Ostwald  was  found  to  hold  for  solu- 
tions of  weakly  dissociated  acids  and  bases;  but  did  not  hold 
for  a  single  strongly  dissociated  compound.  Very  good  con- 
stants were  obtained  for  a  large  number  of  weak  acids  and 

1  See  Author's  Elements  of  Physical  Chemistry,  4th  edition,  p.  406.    (The 
Macmillan  Co.) 

2  Zeit.  phys.  Chem.,  3,  170,  241,  369  (1889).     See  also  Author's  Elements 
of  Physical  Chemistry,  4th  edition,  p.  411.    (The  Macmillan  Co.) 


180  THE  NATURE  OF  SOLUTION 

weak  bases,  varying  the  dilution  of  the  solution  over  a  wide 
range,  and  these  constants  are  very  significant.  Given  the 
" constants"  of  the  acids  and  bases,  we  know  the  relative 
strengths  of  these  substances.  Knowing  their  relative 
strengths,  we  know  exactly  what  these  acids  and  bases  will 
do  chemically  under  any  given  set  of  conditions.  We  know 
with  what  substances  they  will  react,  and  how  rapidly. 
We  know,  further,  the  conditions  of  equilibrium  when  more 
than  one  of  these  compounds  is  involved.  In  a  word, 
knowing  the  "  constant "  of  an  acid  or  a  base,  we  know  its 
physical  chemistry.  We  obtain  the  "constant"  simply  by 
measuring  the  dissociation  of  the  substance  over  a  range 
of  dilutions,  and  this  is  quickly  and  easily  done.  Physical 
chemical  methods  thus  enable  us  to  learn  more  about  acids 
as  acids,  and  bases  as  bases,  in  a  very  short  while,  than  we 
could  learn  in  a  lifetime  through  the  application  of  chemical 
methods  alone.  The  Ostwald  dilution  law  has  a  good 
physical  basis  upon  which  it  rests,  and  from  which  it  is 
deduced  by  the  only  rigid  scientific  method,  viz.,  the  mathe- 
matical. Unfortunately  it  applies  to  only  weak  acids  and 
weak  bases,  but  why  this  limitation  we  do  not  know.  There 
is,  however,  another  expression  which  holds  for  strong  acids 
and  strong  bases,  and  which  does  not  hold  for  the  weakly 
dissociated  compounds.  This  expression,  which  is  purely 
empirical  and  apparently  has  no  physical  basis  at  all,  was 
found  by  Rudolphi.1 


constant. 


hi  which  OL  is  the  percentage  dissociation,  and  v  the  volume 
of  the  solution  or  the  number  of  liters  containing  a  gram- 
molecular  weight  of  the  electrolyte. 

This  expression  seems  to  hold  for  the  strongly  dissociated 
electrolytes  about  as  closely  as  the  Ostwald  law  does  for 
the  weakly  dissociated  compounds.  In  this  expression  we 
find  the  square  root  of  v,  but  we  have  at  present  no  concep- 

1  Zeit.  phys.  Chem.,  17,  385  (1895). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  181 

tion  of  its  meaning.  Other  equations  connecting  dilution 
and  dissociation  were  found  by  Van't  Hoff,1  and  by  Kohl- 
rausch,2  and  by  Krause.3  These  are  again  more  or  less 
empirical  and  need  not  be  discussed  in  detail  hi  the  present 
connection. 

Conductivity  and  Dissociation  at  Elevated  Tempera- 
tures. —  We  have  seen  that  the  conductivities  of  electrolytes 
have  large  temperature  coefficients,  and  that  these  coefficients 
are  positive.  This  means  that  conductivity  increases 
rapidly  with  rise  hi  temperature.  We  can  measure  the 
conductivities  of  aqueous  solutions  in  open  vessels  obviously 
only  up  to  one  hundred  degrees.  To  study  this  property  at 
still  higher  temperatures,  closed  vessels  must  be  used.  A 
convenient  bomb  for  this  purpose  has  been  devised  by  A.  A. 
Noyes  and  co-workers.  A  steel  bomb  lined  with  platinum, 
into  which  properly  insulated  electrodes  had  been  sealed, 
was  employed  for  this  purpose.  The  temperatures  chosen 
were  the  boiling-points  of  certain  liquids  which  could 
readily  be  obtained  reasonably  pure.  They  were,  18°, 
100°,  156°,  218°,  281°,  and  306°,  and  later  work  was  at 
even  higher  temperatures. 

A  few  of  the  results4  obtained  for  a  few  compounds  are 
given  to  show  the  rate  at  which  the  conductivity  increases 
with  rise  hi  temperature. 


Temperature  KC1  NH4C1  K2SO4  HC1 

18°  130.1  130.7  132.8                 376.0 

100°  414.0  415.0  455.0                 850.0 

156°  625.0  628.0  715.0  1085.0 

218°  825.0  841.0  1065.0  1265.0 

281°  1005.0  1460.0 

306°  1120.0  1176.0  1725.0  1424.0 

The  question  arises,  what  is  the  effect  of  rise  hi  tempera- 
ture on  the  dissociation  of  electrolytes?  It  does  not  follow 
that  because  rise  in  temperature  increases  the  conductivity 

1  Zeit.  phys.  Chem.,  18,  300  (1895). 

2  Ibid,  18,  662  (1895). 

3  Zeit.  Elektrochem,  20,  524  (1914). 

4  Taken  from  Carnegie  Institution  of  Washington,  Publication  No.  63, 
335  (1907). 


182  THE  NATURE  OF  SOLUTION 

of  electrolytes,  that  it  also  increases  their  dissociation.  This 
will  be  seen  at  once  if  we  remember  that  dissociation  is  not 
proportional  to  conductivity,  but  is  the  ratio  between  the 
conductivity  at  the  dilution  in  question  and  at  infinite  dilu- 
tion: 


We  have  seen  that  JUB  increases  rapidly  with  rise  hi  tem- 
perature, but  p,m  might  increase  still  more  rapidly;  hi  which 
case  the  dissociation  would  be  less  at  the  higher  temperature. 
The  effect  of  rise  in  temperature  on  dissociation  can  be 
answered  only  by  direct  measurements  of  dissociation  at 
different  temperatures.  This  has  been  done  in  the  same 
experiments  above  referred  to.  Not  only  were  the  values 
of  fj,v  determined  at  the  different  temperatures,  but  the 
effect  of  temperature  on  the  maximum  molecular  conduc- 
tivity was  also  studied. 

The  effect  of  rise  in  temperature  on  dissociation  can  be 
seen  from  a  few  results1  with  a  few  compounds. 

Electrolyte  Concentration  18°  218°  306° 

HC1  0.01  97.1  92.2  82.0 

KC1  0.01  94.2  89.8  81.0 

K2SO4  0.01  87.2  63.0  37.0 

CHaCOOH  0.08  1.50  0.46  0.14 

NH4OH  0.08  1.45  0.47  0.11 

The  results  suffice  to  illustrate  the  general  effect  of  rise 
in  temperature  on  dissociation.  The  higher  the  temperature 
of  the  solution  the  kss  its  dissociation.  The  reason  for  this 
will  appear  in  the  next  chapter,  where  the  relation  between 
the  dissociating  power  of  solvents  and  another  physical 
property  of  liquids  will  be  discussed. 

How  Electrolytes  are  Dissociated  by  Water.  —  We  have 
seen  that  the  characteristic  ion  of  all  acids  is  hydrogen. 
This  means  that  all  acids  dissociate  into  a  hydrogen  ion,  or 
into  hydrogen  ions,  which  carry  the  positive  electricity,  and 
into  something  else  which  carries  the  negative  charge. 

1  Carnegie  Institution  of  Washington,  Publication  No.  63,  340  (1900). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  183 

The  way  in  which  monobasic  acids  break  down  is  very 
simple.  This  is  illustrated  by  hydrochloric  acid, 

HC1=H+  Cl 

The  molecule  is  simply  broken  down  into  one  univalent, 
positive  hydrogen  ion,  and  one  univalent,  negative  chlo- 
rine ion. 

In  the  case  of  dibasic  acids  the  dissociation  is  more 
complicated.  It  takes  place  in  two  stages.  Sulphuric  acid, 
in  fairly  concentrated  solutions,  dissociates  wholly  or  in 
part  as  a  monobasic  acid, 

H2S04  =  H  +  HS04 

When  the  dilution  of  the  solution  is  sufficiently  increased 
we  have  the  second  stage  of  the  dissociation,  which  con- 
sists in  the  dissociation  of  the  HSO4  ion,  thus, 

HS04=H  +  S04 

and  these  two  stages  hi  the  dissociation  of  dibasic  acids 
are  characteristic  of  substances  of  this  class. 

Tribasic  acids  dissociate  hi  three  stages,  as  follows  — 

H3P04=H2P04+H 
H2P04=HP04+H 
HP04=  P04-hH 

Which  of  these  takes  place  depends  upon  the  dilution  of 
the  solution.  In  solutions  of  the  proper  concentrations 
we  may  have  all  three  stages  going  on  simultaneously. 

When  we  turn  to  bases,  similar  phenomena  manifest 
themselves.  A  monacid  base  dissociates  thus, 

NaOH=Na+OH 

Diacid  bases  may  dissociate  hi  two  stages,  thus, 


184  THE  NATURE  OF  SOLUTION 

Ba(OH)2=BaOH+OH 
and  then,  with  increase  in  dilution, 

BaOH=  Ba+OH 

When  we  come  to  salts,  the  problem  of  how  they  dis- 
sociate is  a  very  complex  one,  except  for  the  simple  salts, 
which  dissociate  about  as  we  would  expect  them  to  do, 

KC1=  K+6i 

Ba(N03)2  =  Ba  +  N03  +  NO3 

N+a  +  Na  +  SO4 


These  dissociations  may  go  on  in  stages  with  increase  in 
the  dilution.  But  the  complex  salts  dissociate  in  a  variety 
of  different  ways. 

Take  certain  complex  double  sulphates  —  the  alums  — 
KA1(S04)2;  NH4Cr(SO4)2.  How  do  these  substances  break 
down  in  the  presence  of  water?  What  are  aqueous  solutions 
of  the  alums?  This  problem  was  studied  by  Jones  and  Mac- 
kay.1  They  found  that  in  dilute  solution  the  alums  dissoci- 
ate almost  completely  into  their  constituent  sulphates,  and 
these  are  broken  down  by  the  water  as  if  they  were  alone  in 
separate  solutions.  In  brief,  the  alums  in  dilute  aqueous  so- 
lution are  completely  dissociated  into  their  simplest  ions  thus, 


KA1(S04)2  =  K  +  Al  +  S04  +  S04. 

An  alum  in  dilute  aqueous  solution  is  therefore  not  an 
alum  at  all,  but  is  simply  a  mixture  of  the  ions  from  the 
two  sulphates  of  which  the  alum  in  question  is  composed. 
In  more  concentrated  solutions,  however,  the  alums  first 
break  down  as  salts  of  complex  acids,  thus, 

KA1(S04)2  =  K  +  A1(S04)2, 

and  then  with  increase  in  the  dilution  this  complex  anion 
undergoes  further  dissociation. 

1  Amer.  Chem.  Journ.,  19,  83  (1897). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  185 

The  way  in  which  this  was  proved  may  be  of  some 
interest.  The  freezing-point  lowerings  produced  by  the 
alum  solutions  were  measured.  Also  the  freezing-point 
lowerings  of  the  individual  sulphates,  and  the  sum  of  the 
latter  compared  with  the  former.  If  the  alum  dissociated 
as  a  salt  of  a  complex  acid,  the  sum  of  the  lowerings  of 
the  constituents  would  be  greater  than  that  of  the  alum. 
If  the  latter  dissociated  into  its  simplest  ions,  just  as  the 
constituent  sulphates  would  do,  then  the  sum  of  the  lower- 
ings would  just  equal  the  lowering  produced  by  the  alum 
itself. 

In  all  of  this  work  the  question  of  the  isohydric  or 
non-isohydric  nature  of  the  solutions  was  taken  into 
account,  and  the  effect  of  each  on  the  magnitude  of  the 
dissociation  of  the  other. 

The  term  isohydric  was  first1  applied  to  those  solutions 
which  contain  in  a  given  volume  the  same  number  of  hy- 
drogen ions.  The  term  has  now  become  generic,  and  means 
simply  two  or  more  solutions  which  contain  the  same 
number  of  any  given  kind  of  ion  in  the  same  volume. 
Thus,  two  solutions  which  contain  in  a  given  volume 
the  same  number  of  chlorine  ions  are  termed  isohydric. 

In  attacking  the  above  problem  the  conductivity 
method  was  used  as  well  as  the  freezing-point.  The  con- 
ductivities of  the  individual  solutions  of  the  constituents 
were  measured,  and  then  the  conductivities  of  the  alums, 
and  the  results  compared  exactly  as  their  freezing-point 
lowerings  were  compared,  with  the  same  general  result. 

The  double  chlorides,  bromides,  and  iodides  were  studied 
by  exactly  the  same  methods  that  had  been  used  with 
the  double  sulphates,  with  the  following  result.  The 
double  halides,  as  they  are  termed,  dissociate  as  salts  of 
complex  acids  in  more  dilute  solutions  than  do  the  alums. 
It  requires  greater  dilution  to  break  down  these  com- 
plexes than  in  the  case  of  the  double  sulphates.  Thus,  the 
double  chloride  of  potassium  and  zinc  dissociates  thus, 

1  Wied.  Ann.,  30,  51  (1889). 


186  THE  NATURE  OF  SOLUTION 

K2ZnCl4=K+K+ZnCl4 

and  it  requires  very  great  dilution  to  break  down 
into, 

ZnCU  =  Zn+  C1+  C1+  C1  +  Cl, 

What  has  been  said  of  the  double  chlorides  applies  also 
to  the  double  bromides  and  iodides.  In  concentrated 
solution  they  tend  to  dissociate  as  salts  of  complex  acids, 
and  these  complexes  are  broken  down  only  at  very  great 
dilutions. 

Ways  in  which  Ions  are  Formed.  —  It  has  been  pointed 
out  by  Ostwald  1  that  ions  are  formed  from  molecules  in  a 
number  of  different  ways.  Ions  are  formed  most  fre- 
quently by  the  molecules  simply  breaking  down  into  an 
equivalent  number  of  cations  and  anions.  This  is  the 
way  ions  are  produced  when  acids,  bases,  and  salts  are  dis- 
solved in  water  and  in  other  dissociating  solvents.  Nitric 
acid  in  the  presence  _of  water  is  dissociated  into  the  hydro- 
gen ion  and  the  NO3  anion.  All  acids  are  dissociated 
into  the  hydrogen  ion  or  hydrogen  ions  and  an  anion  the 
nature  of  which  depends  upon  the  acid  in  question. 

As  a  typical  base  take  potassium  hydroxide.  It  dis- 
sociates into  the  hydroxyl  anion  and  the  potassium 
cation  —  all  bases  yielding  hydroxyl  as  the  anion. 

Salts  dissociate  into  cations  and  anions,  the  nature  of 
both  depending  on  the  nature  of  the  salt  in  question. 

The  second  mode  of  ion  formation  pointed  out  by 
Ostwald  is  where  a  metal  like  zinc  displaces  copper  from 
a  solution  of  one  of  its  salts,  or  chlorine  displaces  the 
anion  iodine  from  its  salts.  What  takes  place  in  the 
first  case  is,  the  atom  of  zinc  takes  the  two  positive 
charges  from  the  copper  ion,  liberating  the  atoms  of 
copper;  the  zinc  atom,  becoming  ionic,  passes  into  solu- 
tion. Similarly,  the  chlorine  atom  takes  the  negative 
charge  from  the  iodine  ion,  iodine  being  set  free  and  the 

1  Lehrb.  d.  allg.  Chemie,  II,  786  (1893). 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES  187 

chlorine  becoming  an  anion.  Both  of  these  are  examples 
of  transference  of  an  electrical  charge,  and  this  is  the  key 
to  all  acts  of  substitution  in  chemistry,  whether  organic  or 
inorganic. 

Other  examples  of  this  method  of  ion  formation,  cited 
by  Ostwald,  are  the  replacement  of  gold  from  solutions 
of  auric  chloride  by  hydrogen  gas  under  pressure;  the 
action  of  potassium  on  water,  the  solution  of  metals  in 
general  in  acids,  etc. 

The  third  method  of  ion  formation  cited  by  Ostwald 
is  where  one  neutral  substance  passes  into  cations, 
another  neutral  substance  at  the  same  tune  passing  into 
anions.  The  example  given  is  the  action  of  chlorine  water 
on  gold.  Neither  gold  nor  chlorine  alone  can  become 
ionic,  but  when  brought  into  the  presence  of  one  another, 
the  one,  gold,  takes  positive  charges  and  becomes  cations; 
the  other,  chlorine,  takes  negative  charges  and  becomes 
anions. 

The  fourth  and,  according  to  Ostwald,  the  only  other 
simple  way  in  which  ions  are  formed  is  illustrated  by 
the  action  of  chlorine  on  ferrous  chloride.  The  ferrous 
ion  with  its  two  positive  charges  is  converted  into  the 
ferric  ion  with  three  positive  charges;  the  chlorine  atom 
being  converted  into  the  chlorine  anion.  This  is  oxidation 
hi  that  sense  of  the  term  oxidation  which  means  simply 
raising  the  valence  —  an  unfortunate  use  of  the  term, 
which  should  be  reserved  for  those  chemical  processes 
involving  the  addition  of  oxygen. 

The  same  mode  of  ion  formation  is  illustrated  by  the 
action  of  chlorine  on  a  solution  of  potassium  manganate. 
The  chorine  atom  takes  one  negative  charge  from  the 
ion  MnO4,  converting  it  into  Mn04  which  is  the  anion 
of  permanganates;  one  of  the  potassium  ions  from  the 
manganate  pairing  off  against  the  chlorine  ion  formed  as 
indicated  above. 

These  are  the  simple  ways  in  which  ions  are  formed. 
More  than  one  of  these  may  be  operative  at  the  same  tune 


188  THE  NATURE  OF  SOLUTION 

and  give  rise  to  relatively  complex  processes.  However,  it 
is  not  necessary  to  consider  these  more  complicated  reac- 
tions in  any  detail,  having  hi  mind  the  simple  principles 
underlying  them. 

VELOCITIES  WITH  WHICH  THE  IONS  MOVE 

Faraday's  law  states  that  when  the  electric  current 
is  passed  through  solutions  of  electrolytes,  the  ions  move 
to  the  poles  —  the  cations  to  the  cathode  and  the 
anions  to  the  anode.  The  question  that  arises  in  this  con- 
nection is,  how  rapidly  do  these  ions  move  towards  the 
poles? 

This  involves  two  questions.  First,  what  are  the  rela- 
tive velocities  of  the  different  anions  and  cations,  and 
secondly,  what  are  their  actual  velocities? 

Relative  Velocities  of  Jons.  —  The  determination  of  the 
relative  velocities  with  which  the  ions  move  is  based 
upon  the  changes  in  the  concentrations  of  the  solutions 
around  the  poles,  when  the  current  is  passed  through  the 
solution.  The  change  in  the  concentration  of  the  original 
solution  around  the  cathode,  divided  by  the  amount 
of  the  electrolyte  decomposed,  gives  the  relative  velocity 
of  the  anion.  The  change  in  concentration  around  the 
anode,  divided  by  the  amount  of  the  electrolyte  decom- 
posed, gives  the  relative  velocity  of  the  cation.1  The 
important  point  experimentally  is  to  use  a  form  of 
apparatus2  which  will  allow  a  complete  quantitative 
separation  of  the  solutions  around  the  two  poles,  after 
the  electrolysis  is  completed. 

Apparatus  designed  by  the  Author.  —  A  form  of 
apparatus  which  was  found  to  work  very  satisfactorily, 
is  sketched  in  figure  5.  The  solution  is  placed  in  the 
side  tubes,  which  are  connected  with  a  large  siphon  carry- 
ing a  large  stop-cock.  The  two  electrodes  are  inserted 

1  Amer.  Chem,  Journ.,  32,  409  (1904);  36,  427  (1906). 

2  See  Author's  Elements  of  Physical  Chemistry,  4th  edition,  p.  379  ff. 
(The  Macmillan  Co.) 


AQUEOUS  SOLUTIONS  OF  ELECTROLYTES 


189 


into  the  two  side  tubes,  as  shown  hi  the  drawing.  The 
current  is  passed  until  enough  change  in  concentration 
around  the  poles  is  produced  to  be  measured  with  accuracy, 
and  at  the  same  tune  to  leave  some  unaltered  solution 


hi  the  tops  of  these  tubes.  When  the  current  is  cut  off 
the  stop-cock  is  closed,  the  solutions  removed  from  the 
two  sides  and  analyzed;  the  two  sides  of  the  apparatus 
having  been  calibrated  in  advance. 

To  ascertain  how  much  of  the  electrolyte  has  been 
decomposed  by  the  current,  the  amount  of  current  passed 
is  determined  by  inserting  a  voltmeter  hi  its  path. 
Knowing  the  amount  of  electricity  that  was  passed 


190  THE  NATURE  OF  SOLUTION 

through  the  solution,  from  Faraday's  law  we  can  calculate 
directly  the  amount  of  the  electrolyte  decomposed.  Know- 
ing the  change  in  the  concentration  of  the  electrolyte 
around  each  of  the  poles,  and  the  total  amount  of  the 
electrolyte  decomposed,  we  have  all  the  data  necessary 
for  calculating  the  relative  velocities  of  the  ions. 

Certain  conditions  have  to  be  observed  in  this  work, 
such  as  the  strength  of  the  solution  used,  the  temperature 
at  which  the  work  is  done,  the  strength  of  the  current 
employed,  etc. 

It  should  be  stated  that  in  all  such  work  very  small 
currents  should  be  used.  If  the  current  is  large,  heating 
effects  result  which  tend  to  mix  the  altered  with  the  un- 
altered portion  of  the  solution.  Therefore,  currents  of  only 
a  few  thousandths  of  an  ampere  are  used. 

Effects  of  Concentration  and  Temperature.  —  The  ques- 
tions of  the  concentration  of  the  solution  and  of  the  tem- 
perature, in  their  effects  on  the  relative  velocities  of  the 
ions,  are  of  interest  in  connection  with  certain  work  which 
has  been  done  in  this  laboratory  and  which  will  be  discussed 
in  the  later  chapters  of  this  book. 

It  was  found  that  the  relative  velocities  change  with 
the  concentration  of  the  solution.  The  question  is,  why? 
There  are  at  least  two  reasons.  In  the  first  place  the 
viscosity  of  the  solvent  becomes  less  with  increase  in 
dilution.  This  would  affect  differently  the  velocities  of 
the  two  ions,  if  these  ions  have,  as  they  usually  do,  dif- 
ferent volumes  and  different  masses.  When  we  take  into 
account  the  change  in  hydration  in  aqueous  solutions  with 
change  in  concentration,  as  is  done  in  the  later  chapters  of 
this  work;  and  further,  that  there  is  a  different  change 
in  the  hydration  of  the  cation  and  of  the  anion  with  change 
in  concentration;  we  can  see  how  the  relative  velocities 
of  the  cation  and  the  anion  would  be  changed  with  change 
in  the  concentration  of  the  solution.  The  rule  in  such 
work  is  to  use  a  solution  so  dilute  that,  when  the  dilution  is 
still  further  increased,  the  relative  velocities  remain  constant. 


AQUEOUS   SOLUTIONS   OF   ELECTROLYTES  191 

The  effect  of  temperature  on  the  relative  velocities  of 
the  ions  was,  for  a  long  tune,  a  puzzle.  It  was  early  shown 
by  Loeb  and  Nernst,1  that  the  effect  of  rise  in  tem- 
perature is  to  cause  all  the  ions  to  have  the  same  velocity. 
The  explanation  of  this  rather  surprising  fact  was,  for  a 
long  tune,  not  forthcoming.  The  explanation  today  seems 
to  be  perfectly  simple.  The  effect  of  rise  in  temperature, 
as  we  shall  see  in  later  chapters,  is  to  cause  the  hydrates 
around  the  ions  in  water,  and  the  solvates  around  the 
ions  in  solvents  in  general,  to  become  simpler — the  ten- 
dency of  rise  in  temperature  is  to  de-solvate  the  ions. 
Since  the  different  ions  have  very  different  powers  to  com- 
bine with  the  molecules  of  any  given  solvent,  when  these 
ions  are  even  partially  de-solvated,  they  would  have  more 
nearly  the  same  volumes  and  the  same  masses,  than  when 
they  were  solvated  to  a  greater  extent. 

We  must  further  take  into  account  the  fact,  that  with 
rise  in  temperature  the  viscosities  of  solvents  become  less 
and  less.  Taking  both  of  these  facts  into  account,  we 
can  see  why  rise  in  temperature  would  tend  to  equal- 
ize the  relative  velocities  with  which  the  different  ions 
move. 

Results  of  Measurements  of  the  Relative  Velocities  of 
the  Elementary  Ions.  —  It  was  early  found  that  the 
swiftest  of  all  the  ions  is  hydrogen.  This,  it  will  be  remem- 
bered, is  the  characteristic  ion  of  acids,  is  common  to  all 
acids,  and  always  shows  acidity.  Further,  it  is  one  of  the 
constituents  of  water,  combining  with  the  hydroxyl  ion  to 
form  water. 

Next  to  hydrogen  the  hydroxyl  ion  is  the  swiftest.  Its 
velocity  is  a  little  more  than  half  that  of  the  hydrogen 
ion.  It  should  be  recalled  that  hydroxyl  is  the  charac- 
teristic ion  of  all  bases.  It  is  worthy  of  note  that  the 
characteristic  ions  of  acids  on  the  one  hand,  and  of  bases 
on  the  other,  are  the  swiftest  of  all  the  ions,  and  not  only 
the  swiftest,  but  very  much  the  swiftest,  as  the  following 

1  Zeit.  phys.  Chem.,  2,  962  (1888). 


192  THE  NATURE  OF  SOLUTION 

results  will  show.     These  results  are  taken  from  the  work 
of  Bredig.1 

Element  Atomic  weight  Velocity 

Caesium  132.81  73.6 

Rubidium  85.45  73.5 

Potassium  39.14  70.6 

Sodium  23.00  49.2 

Lithium  6.94  39.8 

Barium  $  137.37  64.0 

Strontium  \  87.63  63.0 

Calcium  *  40.07  62.0 

Magnesium  i  24.32  58.0 

Iodine  126.92  72.0 

Bromine  79.92  73.0 

Chlorine  35.46  70.2 

Fluorine  19.0  50.8 


The  velocity  of  hydrogen  at  25°  is  325,  and  of  hydroxyl 
is  170;  thus,  we  see  that  the  velocity  of  the  hydrogen  ion 
is  several  times  greater  than  that  of  any  other  cation,  and  the 
velocity  of  the  hydroxyl  ion  more  than  twice  that  of  any 
other  anion.  The  relatively  great  velocities  of  the  ions  of 
which  water  is  composed  thus  stand  out  as  among  their 
most  striking  properties. 

Interpretation  of  the  Results.  —  It  is  of  interest  to  note  that 
the  velocities  of  the  alkalies  decrease  as  the  atomic  weights 
decrease.  Lithium  moves  more  slowly  than  sodium,  which 
moves  more  slowly  than  potassium,  rubidium,  and  caesium. 
This  is  exactly  the  opposite  of  what  might  be  expected. 
The  lighter  the  ion,  other  things  being  equal,  the  faster  it 
would  move.  The  case  is  even  more  remarkable,  when  we 
consider  that  the  volume  of  the  lithium  ion  is  less  than 
that  of  sodium;  sodium  less  than  potassium;  potassium 
less  than  rubidium,  and  rubidium  less  than  caesium. 

The  lithium  ion  is  not  only  lighter  than  any  other  ion 
of  the  alkali  group,  but  it  is  also  smaller,*  and  both  of  these 
facts  would  tend  to  make  it  move  more  rapidly  under  a 
given  driving  force.  The  fact  is  it  moves  more  slowly,  and 

1  Zeit.  phys.  Chem.,  13,  242  (1894). 

2  See  Elements  of  Physical  Chemistry,  by  the  author;   4th  edition,  p.  29. 
(The  Macmillan  Co.) 


AQUEOUS   SOLUTIONS   OF    ELECTROLYTES  193 

this  fact  puzzled  physical  chemists  for  some  time.  We 
have  today  a  perfectly  satisfactory  explanation  of  these 
facts. 

Lithium  ions  are  strongly  hydrated l  hi  aqueous  solutions, 
sodium  ions  less  strongly  hydrated,  while  potassium,  ru- 
bidium, and  caesium  ions  are  scarcely  hydrated  at  all. 
What  moves  through  the  solvent  is  the  hydrated  lithium 
ion,  i.e.,  the  lithium  ion  with  its  attached  molecules  of 
water.  This  complex  is  probably  heavier  and  has  a  larger 
mass  than  any  other  alkali  ion,  and  therefore  moves  more 
slowly.  Similarly,  the  hydrated  sodium  ion  has  a  greater 
mass,  and  probably  a  greater  volume,  than  the  much  less 
hydrated  potassium,  rubidium,  and  caesium  ions;  and  there- 
fore moves  more  slowly. 

When  we  come  to  the  alkaline  earths,  we  find  the  same 
general  relations.  Magnesium  ions  are  the  most  strongly 
hydrated,  calcium  and  strontium  next,  and  barium  least, 
and  this  is  the  order  of  their  velocities. 

We  know  at  present  very  little  about  the  hydration  of 
anions.  From  the  above  data  of  relative  velocities  we 
would  be  led  to  conclude  that  fluorine  is  more  hydrated 
than  chlorine,  which  is  more  hydrated  than  bromine,  which, 
in  turn,  is  more  hydrated  than  iodine.  When  the  facts 
are  known,  it  will  be  interesting  to  see  whether  this 
prediction  is  verified. 

The  relation  between  water  of  hydration  in  aqueous 
solutions  of  salts,  and  their  ionic  volumes,  will  be  discussed 
in  one  of  the  later  chapters  of  this  book. 

Bredig2  has  plotted  the  relative  velocities  of  the  ions 
against  the  atomic  weights  as  a  curve,  and  finds  well-defined 
maxima  and  minima  in  the  curve.  These  maxima  and 
minima  are  analogous  to  those  in  the  atomic  volume  curve.3 
The  analogy  is,  indeed,  very  striking.  At  the  maxima 

1  See  Carnegie  Institution  of  Washington,  Publication  No.  60,  31  (1907). 
7  Zeit.  phys.  Chem.,  13,  243  (1894),  also  the  Author's  Elements  of  Physical 
Chemistry,  4th  edition,  p.  387.    (The  Macmillan  Co.) 

»  The  Author's  Elements  of  Physical  Chemistry,  4th  edition,  p.  29. 


194  THE  NATURE  OF  SOLUTION 

of  the  atomic  volume  curve  fall  the  alkali  elements, 
lithium,  sodium,  potassium,  rubidium,  and  caesium.  At 
the  maxima  of  the  ionic  velocity  curve  fall  these  same  ele- 
ments. This  resemblance  between  the  two  curves  mani- 
fests itself  for  many  of  the  elements  other  than  those 
referred  to  above,  but  this  cannot  be  discussed  here  in 
more  detail. 

Actual  Velocities  with  Which  the  Ions  Move.  — The 
question  of  the  relative  velocities  of  the  different  ions  is 
one  thing,  and  that  of  the  actual  velocities  of  the  ions 
under  a  given  driving  force  is  another.  Having  deter- 
mined the  relative  velocities  of  the  different  ions,  and  the 
absolute  velocity  of  any  one  ion,  of  course,  the  absolute 
velocities  of  all  the  ions  would  be  known. 

Our  problem,  then,  is  to  determine  the  absolute  velocity 
of  some  one  ion.  We  shall  see  that  the  fastest  ions  move 
very  slowly.  Therefore,  we  would  naturally  choose  one 
of  the  swiftest  ions  to  measure  the  absolute  velocity,  and 
such  has  been  done.  Indeed,  the  best  measurements  of 
absolute  velocities  of  ions  have  been  made  on  the  very 
swiftest  ion,  hydrogen. 

The  method  we  owe  to  the  English  physicist,  Lodge.1 
We  mean  by  the  velocity  of  an  ion  the  same  that  we 
mean  by  the  velocity  of  any  moving  body  —  the  dis- 
tance traveled  hi  a  given  time.  How  far  does  the  hy- 
drogen ion  travel  under  a  given  driving  force,  in  unit 
time?  The  apparatus  and  method  used  by  Lodge  are  the 
following. 

Two  beakers  were  connected  with  a  graduated  siphon, 
and  into  each  beaker  a  platinum  electrode  was  plunged. 
Sulphuric  acid  was  poured  into  the  two  beakers,  and  the 
connecting  siphon  was  filled  with  jelly  containing  sodium 
chloride  to  which  just  enough  alkali  was  added  to  show  the 
alkaline  reaction  with  phenolphthalein  which  was  also 
dissolved  in  the  jelly.  In  Lodge's  own  words,  "To  detect 
the  motion  of  hydrogen,  Mr.  Robinson  devised  the  follow- 

1  Brit.  Ass.  Report,  1886,  p.  393. 


AQUEOUS   SOLUTIONS   OF   ELECTROLYTES  195 

ing  arrangement: — We  happened  to  have  been  using 
phenolphthalein  as  a  detector  of  alkali  in  some  other 
quite  distinct  experiment,  and  so  it  was  a  handy  substance. 
The  jelly  tube  contains  a  little  phenolphthalein  and  a  trace 
of  common  salt,  just  made  alkaline  enough  with  soda  to 
bring  out  the  color.  The  solution  in  the  anode  vessel  is 
H2S04;  in  the  cathode  vessel  the  same,  or  sometimes 
CuS04. 

"The  result  is  that  S04  travels  one  way,  and  H  the 
other.  As  the  H  travels,  it  liberates  HC1,  and  decolorizes 
the  solution  by  forming  neutral  Na^SC^.  The  velocity  of 
hydrogen,  for  40  volts  applied  to  a  40  centimeter  tube, 
came  out  from  the  very  first  observation  thus  made  as 
0.0029  centimeters  per  second.  Kohlrausch's  theoretical 
number,  deduced  from  conductivity  and  migration  data,  is 
0.003.  Later  experiments  gave  respectively  .0026  and  .0024. 
S04  seems  to  travel  at  about  one- third  this  speed." 

So  much  for  Lodge's  method  and  for  the  results  obtained 
by  him.  For  a  potential  drop  of  one  volt  per  centimeter, 
along  the  tube,  the  hydrogen  ion,  the  swiftest  of  them 
all,  moves  with  a  velocity  of  about  three-thousandths  of 
•  a  centimeter  per  second;  and  the  other  ions  with  velocities 
which  bear  the  relation  to  this,  of  the  relative  velocity  of 
the  ion  in  question  to  the  relative  velocity  of  the  hydrogen 
ion.  Thus,  chlorine  and  potassium,  which  have  practically 
the  same  velocities,  have  under  the  conditions  of  the 
Lodge  experiment,  an  absolute  velocity  of  -B^  of  0.003  cm. 
per  second;  and,  in  a  similar  manner,  the  absolute  veloci- 
ties of  any  other  ion  whose  velocity  relative  to  hydrogen 
is  known,  can  be  calculated. 

We  thus  see  that  the  swiftest  ions  move  very  slowly 
indeed,  even  under  a  strong  driving  force. 

Velocities  of  Ions  and  of  Gaseous  Molecules.  —  It  is 
interesting  to  compare  the  velocity  of  the  hydrogen  ion 
in  solution  under  an  electrical  force  of  a  potential  drop 
of  a  volt  per  centimeter,  with  the  velocity  of  the  gaseous 
hydrogen  molecule  under  normal  conditions.  We  have  seen 


196  THE  NATURE  OF  SOLUTION 

that  under  the  conditions  of  the  Lodge  experiment,  the 
hydrogen  ion  moves  only  0.003  cm.  per  second.  The 
hydrogen  molecule  in  hydrogen  gas,  under  normal  conditions 
of  temperature  and  pressure,  moves  with  a  velocity  of 
over  a  mile  a  second  —  the  velocity  being  of  an  entirely 
different  order  of  magnitude  from  that  at  which  the  hydro- 
gen ion  travels. 

This  comparison  is  of  interest  in  connection  with  the 
relations  between  the  gas-pressure  of  a  gas  and  the  os- 
motic pressure  of  a  dissolved  substance,  which  have  already 
been  pointed  out  (p.  72).  We  know  the  cause  of  gas- 
pressure.  It  is  explained  by  the  kinetic  theory.  It  is 
due  to  the  gas-particles  striking  against  the  walls  of  the 
containing  vessel.  The  attempt  was  made  by  Van't  Hoff1 
to  explain  osmotic  pressure  in  a  similar  kinetic  way.  The 
attempt,  however,  was  not  successful;  and  no  other  ki- 
netic theory  of  osmotic  pressure  thus  far  proposed  can  be 
regarded  as  explaining  osmotic  pressure  at  all  satisfactorily. 

ELECTROMOTIVE  FORCE  OF  PRIMARY  CELLS 

One  of  the  most  important  scientific  applications  of  the 
relations  between  solutions  and  gases  and  of  the  theory  of 
electrolytic  dissociation,  is  to  the  problem  of  the  electro- 
motive force  of  primary  cells. 

The  primary  cell  was  discovered  by  Galvani  and  con- 
structed by  Volta  more  than  one  hundred  years  before  its 
action  was  understood.  A  typical  primary  cell  consists 
of  two  metals,  each  surrounded  by  one  of  its  own  salts; 
the  metals  being  connected  with  each  other,  and  the  two 
solutions  connected  through  a  siphon  filled  with  the  one 
or  the  other  solution. 

What  is  the  source  of  the  electromotive  force  in  such  an 
element,  or  what  are  the  sources  of  the  electromotive 
force?  There  are  several  possibilities,  and  widely  different 
views  have  been  held  concerning  the  action  of  the  primary 
cell.  Some  have  supposed  that  the  chief  sources  of  the 

1  Zeit.  phys.  Chem.,  1,  481  (1887). 


AQUEOUS   SOLUTIONS   OF   ELECTROLYTES  197 

electromotive  force  were  at  the  surfaces  of  contact  of  the 
electrodes  with  the  electrolytes.  Others,  that  the  main 
action  of  the  primary  cell  was  at  the  contact  of  the  two 
electrolytes  with  one  another;  and  still  others,  that  the 
source  of  most  of  the  electrical  energy  was  at  the  contact 
of  the  two  metals  with  one  another. 

The  problem  of  the  primary  cell  as  a  unit  —  as  a  ma- 
chine —  was  dealt  with  by  Willard  Gibbs  hi  this  country,  and 
by  Helmholtz  hi  Germany,  on  thermodynamical  grounds. 

The  simplest  form  of  primary  cell  consists  of  two  elec- 
trodes of  same  metal  plunged  into  solutions  of  a  salt  of 
the  metal,  the  two  solutions  around  the  two  electrodes 
having  different  concentrations. 

Gibbs  and  Helmholtz  were  able  to  establish  a  relation 
between  the  electrical  energy  which  appears  in  such  a  cell, 
and  the  heat  energy  required  to  produce  equality  of  con- 
centration on  the  two  sides  by  distilling  water  from  the 
more  dilute  into  the  more  concentrated  solution  until  the 
two  concentrations  were  equal. 

This  was  an  important  step  hi  the  solution  of  the 
problem  of  the  primary  cell,  but  it  dealt  with  the  cell  as 
•a  whole.  It  did  not  go  into  the  cell,  as  it  were,  and  show 
how  it  acted.  It  did  not  analyze  the  cell  and  show  how 
much  of  the  electromotive  force  came  from  one  source  and 
how  much  from  another.  This  was  done  by  the  applica- 
tion of  what  was  learned  about  solutions  by  Van't  Hoff 
and  Arrhenius;  by  the  application  of  the  gas  laws  to  solu- 
tion and  the  dissociation  theory  to  the  problem  in  hand. 

This  application  we  owe  to  Nernst. 

Solution-Tension  of  the  Metals.  —  Nernst  introduced 
a  conception,  new  at  the  tune,  which  he  called  the  elec- 
trolytic solution-tension  of  the  metals.  In  a  word,  it  is 
this.  When  a  bar  of  metal  is  immersed  hi  a  solution  of 
one  of  its  own  salts  (or  into  other  solutions  or  solvents), 
there  is  at  its  surface  a  tension  or  pressure  which  tends 
to  drive  the  metal  atoms  hi  the  form  of  ions  off  the  bar 
into  the  solution. 


198  THE  NATURE  OF  SOLUTION 

Opposing  the  solution-tension  of  the  metal  is  the 
osmotic  pressure  of  the  ions  in  the  solution.  In  some 
forms  of  cells  it  is  the  osmotic  pressure  of  the  cation  which 
is  operative;  in  other  forms,  the  anion;  and  in  still  other 
forms,  the  osmotic  pressure  of  both  ions  comes  into  play. 
Where  the  osmotic  pressure  of  the  cation  is  the  condi- 
tioning factor,  we  have  this  opposing  the  solution-tension 
of  the  metal,  and  what  will  occur  depends  upon  the  rela- 
tive magnitudes  of  these  two  forces.  Where  the  solution- 
tension  of  the  metal  is  greater  than  the  osmotic  pressure 
of  the  cations,  metal  atoms  separate  from  the  bar,  take  a 
positive  charge  from  it  and  pass  into  solution  as  cations, 
leaving  the  bar  charged  negatively. 

When  the  osmotic  pressure  of  the  cations  in  the  solu- 
tion is  greater,  metal  ions  separate  from  the  solution  on  to 
the  bar,  give  up  their  positive  charges  to  the  bar,  and  the 
electrode  is  thus  charged  positively.  This  is  the  key  to  the 
action  of  the  primary  cell.  The  question  that  remains  is 
this.  Is  the  assumption  of  the  solution-tension  of  the 
metals  founded  on  fact?  Is  there  any  such  force  at  the 
surface  of  contact  of  a  metal  and  a  solution  say  of  one  of 
its  salts? 

Experimental  Proof  of  the  Existence  of  Solution-Ten- 
sion.—  The  fundamental  point  in  connection  with  these 
deductions  of  Nernst  is  the  existence  of  solution- tension. 
Can  it  be  demonstrated?  At  first  it  was  a  pure  assump- 
tion. Later,  the  existence  of  solution-tension  was  demon- 
strated by  Palmaer,1  in  the  following  manner. 

A  vessel  was  filled  with  a  solution  of  mercurous  nitrate. 
Very  finely  divided  mercury  was  allowed  to  rain  down 
through  the  solution  in  the  form  of  a  mercury  mist.  Mer- 
cury is  an  element  with  a  very  low  solution-tension,  as  we 
shall  see.  This  means  that  when  mercury  comes  in  con- 
tact with  a  solution  of  one  of  its  own  salts,  mercury  ions 
separate  from  the  solution  on  to  the  drop,  charging  it 
positively.  As  each  positively  charged  droplet  of  mercury 

1  Zeit.  phys.  Chem.,  25,  265  (1898);  28,  257  (1899). 


AQUEOUS   SOLUTIONS   OF   ELECTROLYTES  199 

falls  through  the  solution,  it  draws  to  it  electrostatically 
the  NO3  anions  in  the  solution,  and  carries  these  down 
with  it  to  the  bottom  of  the  vessel.  At  the  bottom  of  the 
beaker  is  placed  some  mercury.  When  the  positively 
charged  falling  drops  come  hi  contact  with  the  mercury, 
the  latter  is  also  charged  positively;  the  N03  anions  which 
were  dragged  down  through  the  solution  by  the  drop 
being  set  free  at  the  surface  of  the  mercury.  The 
mercury  at  the  bottom  of  the  vessel  now  being  charged 
positively,  will  throw  mercury  ions  back  again  into 
the  solution,  _and  these  cations  will  pair  themselves  off 
against  the  NO3  anions  which  now  exist  hi  large  numbers 
just  above  the  surface  of  the  metal.  The  mercurous 
nitrate  should  thus  become  more  concentrated  just  above 
the  surface  of  the  mercury,  than  hi  other  parts  of 
the  solution. 

This  conclusion  is  based  solely  on  the  assumption  of 
the  existence  of  solution-tension  of  the  metals.  The  facts 
confirm  the  conclusion.  When  the  mercury  is  allowed  to 
rain  down  hi  the  form  of  a  very  fine  mist,  the  mercurous 
nitrate  just  above  the  metallic  mercury  becomes  as  much  as 
forty  per  cent  more  concentrated  than  the  remainder  of 
the  solution. 

By  simply  allowing  finely  divided  mercury  to  rain  down 
through  a  homogeneous  solution  of  mercurous  nitrate,  and 
thus  to  change  the  concentration  of  the  solution  appreciably 
is  a  most  surprising  result.  It  could  never  have  been  sus- 
pected until  we  had  the  conception  of  the  solution-tension 
of  metals,  and  its  existence  is  strong  confirmation  of  the 
correctness  of  the  assumption  made  by  Nernst. 

Values  of  the  Solution-Tensions  of  Certain  Metals. — 
It  would  lead  us  too  far  to  discuss  hi  any  detail,  in  the 
present  connection,  the  method  of  determining  the  solution- 
tension  of  the  metals.  For  this  reference  must  be  had  to 
some  text-book1  on  Physical  Chemistry.  The  results  are 

1  See  Author's  Elements  of  Physical  Chemistry,  4th  edition,  p.  487.  (The 
Macmillan  Co.) 


200  THE  NATURE  OF  SOLUTION 

of  such  importance  and  bear  on  so  many  problems  that  for 
some  of  the  more  common  elements  they  are  given  hi  the 
following  table.1 

Atmospheres 

Magnesium  1044 

Zinc  1018 

Aluminium  1013 

Cadmium  107 

Iron  104 

Lead  10~3 

Mercury  10~16 

Silver  KT17 

Copper  10-20 

This  tension-series,  as  it  is  called,  tells  us  just  which 
metals  will  precipitate  other  metals  from  solutions  of  their 
salts,  when  the  metal  is  plunged  into  the  solution.  A 
metal  in  this  tension-series  will,  in  general,  precipitate 
from  solutions  of  its  salts  any  metal  which  is  lower  down  hi 
the  tension-series,  and  be  precipitated  from  its  own  salts  by 
a  metal  higher  in  the  series.  To  secure  this  result,  how- 
ever, the  metals  should  be  sufficiently  widely  removed  from 
one  another  in  the  series — there  should  be  a  sufficient 
difference  in  their  solution-tensions.  The  best  results  are 
secured  when  a  metal  near  the  top  of  the  series,  like  zinc, 
is  plunged  into  the  solution  of  a  metal  near  the  bottom  of 
the  series,  like  copper. 

We  are  impressed  by  the  enormous  differences  in  the 
solution-tensions  of  the  different  metals.  Thus,  mag- 
nesium has  a  solution-tension  of  1044  atmospheres,  and 
copper,  a  solution-tension  of  only  10~20  atmospheres  —  a 
variation  almost  from  infinite  to  infinitesimal. 

At  first  thought,  such  extreme  values  are  apt  to 
impress  one  unfavorably.  These  metals  show  a  general 
resemblance  in  other  physical  properties.  Marked  differ- 
ences, to  be  sure,  manifest  themselves,  but  no  differences 
of  anything  like  the  order  of  magnitude  shown  by  the 
solution-tensions.  Yet,  when  these  experimentally  found 
values  are  used  in  the  equations  for  calculating  the  elec- 
tromotive forces  of  primary  cells  containing  these  metals 

»  Ostwald:  Lehrb.  d.  allg.  Chem.,  II,  948  (1893). 


AQUEOUS   SOLUTIONS   OF   ELECTROLYTES  201 

as  electrodes,  the  calculated  values  are  in  agreement  with 
those  found  experimentally  to  within  the  limit  of  error  of 
experiment,  and  this  is  the  real  test  of  their  worth. 

We  must  accept,  then,  at  least  tentatively,  the  above 
values  for  the  solution-tensions  of  the  metals  as  correspond- 
ing to  the  facts  of  nature. 

An  interesting  application  of  the  solution-tension  of  the 
metals  has  been  made  by  Ostwald1  to  explain  why  pure 
metals  such  as  zinc  do  not  dissolve  hi  acids  and  why  they 
do  dissolve  when  platinum  is  brought  hi  contact  with  them. 
He  showed  that  zinc,  being  a  metal  with  a  high  solution- 
tension,  does  not  dissolve  because  the  hydrogen  ions  of 
the  acid  cannot  give  up  their  charges  to  it  and  escape  as 
hydrogen  gas.  When  the  zinc  is  attached  to  platinum  —  a 
metal  with  a  low  solution-tension  —  the  hydrogen  ions  give 
up  their  charges  to  the  platinum  and  escape  as  hydrogen 
gas,  the  zinc  passing  into  solution. 

Ostwald  carried  out  this  experiment  hi  such  a  way  as  to 
illustrate  also  chemical  action  without  mechanical  contact. 
It  would  lead  us  too  far  to  discuss  this  hi  detail  here.2 
Suffice  it  to  say  that  he  showed  that  electrical  contact 
rather  than  mechanical  contact  is  essential  in  order  that 
things  may  react  chemically;  mechanical  contact  being  hi 
general  the  simplest  means  of  establishing  electrical  contact. 

1  Zeit.  phys.  Chem.,  9,  540  (1892). 

s  See  Elements  of  Physical  Chemistry,  by  the  Author,  4th  edition,  p.  489. 
(The  Macmillan  Co.) 


CHAPTER  XI 

SOLUTIONS  IN  NONAQUEOUS  AND   IN  MIXED   SOLVENTS 

Solvent  Power  of  Liquids.  —  The  power  of  liquids  to 
dissolve  substances  varies  greatly  from  liquid  to  liquid. 
Of  all  known  liquids  water  has  the  greatest  solvent  power, 
but  we  must  not  conclude  that  water  has  a  monopoly, 
or  anything  approaching  it,  as  a  solvent.  All  liquids 
have  some  solvent  power.  Some  liquids  dissolve  a  large 
number  of  substances,  and  certain  substances  hi  large 
quantities.  Thus,  the  alcohols,  simple,  or  complex  like 
glycerol,  have  very  marked  solvent  power.  The  simple 
alcohols  dissolve  in  large  quantities  many  fats,  oils,  resins, 
and  the  like,  which  are  either  practically  insoluble  in 
water,  or  dissolve  to  only  a  very  slight  extent.  Glycerol 
is  an  excellent  solvent,  especially  at  more  elevated  tempera- 
tures. Acetone  dissolves  a  great  variety  of  things,  and 
the  solutions  in  acetone  have  abnormal,  and  therefore 
very  interesting  properties;  acetone  having  the  power 
of  polymerizing  a  great  variety  of  substances  when  dis- 
solved in  it. 

We  also  find  good  solvents  among  the  inorganic  liquids. 
Thus,  liquid  ammonia  has  very  marked  solvent  power; 
liquid  hydrocyanic  acid  is  a  good  solvent  for  quite  a  variety 
of  substances,  and  there  are  many  other  examples. 

The  Relative  Powers  of  Different  Liquids  to  Dis- 
sociate Electrolytes.  —  Let  us  consider  first  certain  com- 
pounds of  carbon.  Take  the  simplest  alcohols.  The 
dissociating  powers  of  these  substances  can  now  be  meas- 
ured by  the  improved  conductivity  method,  and  the  power 
of  methyl  and  ethyl  alcohol  to  break  molecules  down  into 
ions  can  also  be  measured  by  the  improved  boiling-point 


NONAQUEOUS  AND  MIXED  SOLVENTS  203 

method,1  as  has  already  (p.  118)  been  seen.  Methyl 
alcohol  has  from  one-half  to  one-fourth  the  dissociating 
power  of  water;  ethyl  alcohol  from  one-third  to  one-fifth 
the  dissociating  power  of  water.  The  higher  members  of 
this  series  of  alcohols,  i.e.,  those  containing  a  larger  num- 
ber of  carbon  atoms,  have  kss  and  kss  dissociating  power 
the  more  complex  the  alcohol,  and  this  is  true  in  general 
of  members  of  any  homologous  series  of  compounds  of 
carbon.  The  more  complex  the  compound,  therefore, 
the  higher  it  stands  hi  the  series,  the  less  its  power  to 
break  down  molecules  dissolved  in  it,  into  charged  parts 
or  ions. 

For  our  knowledge  of  the  dissociating  powers  of  methyl 
and  ethyl  alcohols,  we  are  indebted  especially  to  the  work 
of  Carrara,2  Zelinsky  and  Krapiwin,3  Vollmer,4  Kablukoff , 5 
and  Fitzpatrick.6  Considerable  work  hi  the  alcohols  as 
solvents  has  also  been  done  hi  this  laboratory,  as  will 
appear  especially  hi  the  later  sections  of  this  chapter, 
which  deal  with  investigations  hi  mixed  solvents. 

The  work  hi  the  higher  alcohols  of  this  series  we  owe 
especially  to  Schlamp,7  Carrara,8  and  Kablukoff.9 

Schlamp  showed  that  propyl  alcohol  has  somewhat 
less  than  one-half  the  dissociating  power  of  ethyl  alcohol; 
thus  illustrating  the  fact  mentioned  above,  that  the  more 
complex  the  member  of  the  homologous  series  the  less 
its  dissociating  power. 

The  work  of  Kablukoff  in  isoamyl  alcohol  brought  out 
a  fact  of  interest  and,  no  doubt,  of  importance,  the  mean- 
ing of  which,  however,  is  at  present  not  understood. 
Hydrochloric  acid  dissolved  in  this  solvent  gave  a  molec- 
ular conductivity  which  decreased  with  increase  in  dilution 
(see  p.  177).  It  will  be  recalled  that  this  is  exactly  the 


Zeit.  phys.  Chem.,  31, 114  (1899).  •  Phil.  Mag.,  24,  378  (1887). 

Gazz.  Chim.  ital,  26, 1, 119  (1896).  7  Zeit.  phys.  Chem.,  14,  272  (1894). 

Zeit.  phys.  Chem.,  21,  35  (1896).  8  Gazz.  Chim.  ital,  27, 1, 221  (1897). 

Wied.  Ann.,  62,  328  (1894).  9  Zeit.  phys.  Chem.,  4,  432  (1889). 
Zeit.  phys.  Chem.,  4,  429  (1889). 


204  THE  NATURE  OF  SOLUTION 

reverse  of  the  general  effect  of  dilution  on  the  molecular 
conductivity  of  solutions  of  electrolytes,  the  more  dilute 
the  solution  the  greater  the  molecular  conductivity. 

Other  Compounds  of  Carbon.  —  A  similar  relation  was 
found  by  Kablukoff1  in  ether  as  a  solvent.  The  molec- 
ular conductivity  of  hydrochloric  acid  in  this  solvent  also 
decreases  with  the  dilution.  These  same  solutions  were 
found  to  show  less  conductivity  the  higher  the  temperature 
to  which  they  were  heated  —  in  a  word,  negative  tem- 
perature coefficients  of  conductivity.  This  is  true  of 
metals  in  general,  as  we  have  seen,  but  is  directly  the 
opposite  of  what  we  find  for  nearly  all  solutions.  The 
hydrocarbons,  such  as  benzene,  etc.,  have  very  little  dis- 
sociating power.  The  same  is  true,  in  general,  of  the 
ethers,  aldehydes,  esters,  and  substitution  products  of  these 
so-called  neutral  organic  compounds. 

Acetone  has  very  pronounced  dissociating  power;  indeed 
about  one-fourth  to  one-fifth  that  of  water.  As  already 
stated,  solutions  hi  acetone  present  a  number  of  abnormal 
properties.  Some  of  the  molecules  are  largely  associated 
and  others  are  simultaneously  broken  down  into  ions. 
The  best  conductivity  work  in  acetone  we  owe  to 
St.  v.  Laszczynski,2  Dutoit  and  Aston,3  and  Dutoit  and 
Friderich.4  Some  work  has  been  done  in  the  more  com- 
plex ketones,  with  the  same  general  result,  that  the  higher 
members  of  a  homologous  series  have  less  dissociating 
power  than  the  lower. 

The  conductivity  of  certain  salts  in  pyridine  was  stud- 
ied by  Werner,5  but  most  of  the  work  on  the  dissociating 
power  of  this  solvent  we  owe  to  St.  v.  Laszczynski  and  St. 
v.  Gorski.6  While  the  dissociating  power  of  this  solvent 
is  not  yet  accurately  known,  still  it  is  known  to  have  very 
considerable  power  of  breaking  molecules  down  into  ions. 

There  is  one  organic  acid  —  formic  acid  —  which  was 

1  Zeit.  phys.  Chem.,  4,  431  (1889).     4  Bull  Soc.  Chim.  [3],  19,  321  (1898). 

2  Zeit.  Elektrochem.,  2,  55  (1895).      6  Zeit.  anorg.  Chem.,  16,  I,  123  (1897). 

3  Campt.  rend.,  125,  240  (1897).         6  Zeit.  Elektrochem.,  4,  290  (1897). 


NONAQUEOUS  AND  MIXED  SOLVENTS  205 

shown  by  the  work  of  Zanniovich-Tessarin,1  to  have  very 
great  dissociating  power.  Indeed,  the  dissociating  power 
of  this  solvent  is  of  the  same  order  of  magnitude  as  that 
of  water  itself. 

The  dissociating  power  of  acetic  acid  was  shown  by 
Jones2  to  be  very  much  less  than  that  of  formic  acid. 

The  dissociating  power  of  glycerol  as  a  solvent  has  been 
studied  at  some  length  hi  the  author's  laboratory;  but  the 
results  obtained  will  be  discussed  in  a  later  section  of  this 
chapter,  dealing  primarily  with  mixed  solvents. 

The  one  to  whom  we  probably  owe  more  than  to  any 
other  for  our  knowledge  of  the  dissociating  power  of  a 
large  number  of  organic  and  inorganic  solvents  is  Paul 
Walden3  of  Riga.  He  has  studied  the  conductivities  of 
electrolytes  when  dissolved  hi  many  unusual  solvents,  and 
has  thrown  light  on  the  dissociating  powers  of  many  liquids. 
It  would  lead  us  too  far  to  discuss  these  elaborate  investiga- 
tions here  hi  any  detail.  A  few  points  established  by  him 
are,  however,  of  such  interest  and  importance  that  they 
must  be  considered. 

Solvents  Other  Than  Carbon  Compounds.  —  Walden 
showed  that  among  the  inorganic  solvents  some  have  very 
marked  dissociating  power,  while  others  have  scarcely  any 
dissociating  power  at  all.  For  example,4  arsenic  and  anti- 
mony trichlorides  and  phosphorus  oxychloride  have  marked 
dissociating  power;  while  the  trichloride  and  tribromide 
of  phosphorus,  the  pentachloride  of  antimony,  and  stannic 
chloride  have  very  little  power  to  break  molecules  down 
into  ions.  Arsenic  tribromide5  also  has  considerable  dis- 
sociating power.  It  is  interesting  to  note  the  difference 
between  the  various  halogen  compounds  of  members  of  the 
phosphorus  group,  with  respect  to  their  power  to  break 
down  molecules  into  ions. 

1  Zeit.  phys.  Chem.,  19,  251  (1896). 

2  Amer.  Chem.  J&urn.,  16,  13  (1894). 

3  Zeit.  phys.  Chem.,  46,  103  (1903);   64,  129  (1906);   65,  207,  281,  683 
(1906). 

4  Zeit.  anorg.  Chem.,  25,  209  )1900).  6  Ibid.,  29,  371  (1902). 


206  THE  NATURE  OF  SOLUTION 

Abnormal  Electrolytes.  —  One  of  the  most  striking 
relations  brought  out  by  Walden1  deals  with  the  so-called 
abnormal  electrolytes.  He  found  that  when  the  halogens, 
phosphorus,  arsenic,  antimony,  sulphur,  and  certain  other 
substances  were  dissolved  in  such  solvents  as  liquid  sul- 
phur dioxide  and  arsenic  chloride,  the  solutions  showed 
marked  conductivity.  How  was  this  possible?  The  dis- 
solved substances  are  elements,  not  electrolytes  at  all 
in  the  proper  sense  of  the  term.  What  kinds  of  ions  do 
they  yield?  In  order  to  conduct  the  current  these  sub- 
stances must  break  down  into  both  positively  charged  and 
negatively  charged  parts.  We  must  have  both  cations  and 
anions  formed  by  these  elementary  substances.  This  is 
obviously  different  from  the  ordinary  electrolytic  dissocia- 
tion of  acids,  bases  and  salts.  Walden  termed  such  sub- 
stances hi  such  solvents,  abnormal  electrolytes. 

If  we  examine  the  work  of  Thomson2  on  the  electrolysis 
of  hydrogen  gas,  in  which  he  showed  that  the  molecule 
is  made  up  of  both  -positive  and  negative  constituents,  we 
will  see  a  certain  analogy  between  what  takes  place  in  the 
hydrogen  when  the  current  is  passed  through  it,  and  what 
takes  place  when  the  above-named  elements  are  dissolved 
in  the  different  solvents.  The  hydrogen  molecule  is 
broken  down  into  a  positive  and  into  a  negative  constit- 
uent, as  is  shown  by  the  different  spectra  on  the  two 
sides  of  the  metal  septum. 

It  seems  probable  that  the  above  elementary  sub- 
tances  are  broken  down  by  certain  solvents  into  both 
positive  and  negative  constituents  or  ions.  Otherwise  it 
would  be  difficult  to  see  how  such  solutions  could  carry  a 
current,  and  show  conductivity  as  it  is  ordinarily  measured. 

Hydrocyanic  Acid. —  Centnerszwer3  has  shown  that  the 
dissociating  power  of  liquid  hydrocyanic  acid,  itself  a  com- 
pound of  carbon,  is  greater  even  than  that  of  water.  This 
is  very  important  as  we  shall  see  in  connection  with  the 

1  Zeit.  phys.  Chem.,  43, 385  (1903).          •  Zeit.  phys.  Chem.,  39, 217  (1902). 

2  Nature,  62,  451  (1895). 


NONAQUEOUS  AND  MIXED  SOLVENTS  207 

relation  between  the  dissociating  powers  of  solvents  and 
their  other  physical  properties. 

While  liquid  hydrocyanic  acid  has  this  remarkably  high 
dissociating  power,  Centnerszwer  has  found  that  liquid 
cyanogen  has  very  small  dissociating  power  indeed.  This 
is  rather  surprising  when  we  consider  that  the  only  dif- 
ference in  chemical  composition  between  these  two  com- 
pounds is  an  atom  of  hydrogen. 

Hydrogen  Dioxide.  —  Hydrogen  dioxide  has  been  shown 
by  Jones1  and  his  co-workers  Barnes  and  Hyde,  to  have  a 
higher  dissociating  power  than  water.  This  is  not  sur- 
prising when  we  consider  the  close  relation  chemically 
between  hydrogen  dioxide  and  water.  The  former  is 
simply  oxidized  water.  This  fact  is  also  of  interest  in 
connection  with  the  relation  just  referred  to  between 
the  dissociating  power  of  liquids  and  certain  of  their 
other  physical  properties. 

Dissociating  Power  of  Liquid  Ammonia.  —  Cady  in 
18972  discovered  that  solutions  of  salts  in  liquid  ammonia 
are  good  conductors  of  the  current.  This  led  Franklin 
and  Kraus3  to  take  up  an  extensive  study  of  the  con- 
ductivity of  salts  dissolved  hi  this  solvent.  They  found 
that  such  solutions  conducted  far  better  than  aqueous 
solutions  of  the  same  salts  at  the  same  concentrations. 
They  at  first  concluded  that  such  solutions  are  more  dis- 
sociated than  aqueous  solutions  of  the  same  salts  having 
the  same  concentrations.  This  conclusion  was  later 
somewhat  modified.  The  conductivity  of  a  solution  is  a 
function  not  only  of  the  number  of  ions  present,  but  also 
of  the  velocities  with  which  they  move.  The  velocities 
of  the  ions  in  liquid  ammonia  were  shown  to  be  much 
larger  than  in  water,  under  the  same  conditions,  as  would 
be  expected  from  the  small  viscosity  of  liquid  ammonia. 

1  Amer.  Chem.  Journ.,  27,  22  (1902). 

2  Journ.  Phys.  Chem.,  1,  707  (1897). 

8  Amer.  Chem.  Journ.,  20,  820,  836  (1898);  21,  8  (1899);  23,  277  (1900); 
24,  83  (1900);  28,  83  (1902);  Journ.  Amer.  Chem.  Soc.,  26,  499  (1904);  27, 
191  (1905);  29,  1389  (1907). 


208  THE  NATURE  OF  SOLUTION 

When  this  was  taken  into  account,  it  was  found  that  salts 
dissolved  hi  liquid  ammonia  are  dissociated  to  about  one- 
fourth  the  extent  that  they  are  hi  water  under  the  same 
conditions. 

Lewis1  carried  out  an  investigation  hi  iodine  as  the 
solvent.  He  found  that  hi  dilute  solutions  the  conductiv- 
ity of  potassium  iodide  dissolved  in  this  solvent  increases 
rapidly  with  concentration  up  to  a  maximum,  after  which  it 
falls  off.  This  would  lead  one  to  suspect  that  there  is  some 
compound  formed  between  the  solvent  and  the  dissolved 
substance. 

In  connection  with  the  work  in  nonaqueous  solvents,  we 
should  mention  that  of  Archibald  and  Mclntosh,  who 
used  as  solvents  the  liquefied  halogen  acids,  hydrochloric, 
hydrobromic,  and  hydriodic,  and  also  liquid  hydrogen 
sulphide.  Organic  compounds  dissolved  hi  these  solvents 
showed  an  increase  in  the  molecular  conductivity  with 
increase  in  the  concentration  of  the  solution.  They  pointed 
out  that  there  are  a  fairly  large  number  of  examples 
illustrating  this  same  condition,  the  number  being  much 
larger  than  any  one  supposed  before  they  were  thus  col- 
lected. 

Davis  and  Putnam,2  working  in  this  laboratory,  have 
made  a  fairly  extensive  study  of  the  conductivity  of  solu- 
tions of  salts  in  formamide  as  the  solvent.  They  have 
found  that  formamide  has  a  greater  dissociating  power 
than  water;  salts  being  completely  dissociated  hi  it  at 
lower  dilutions  than  hi  water.  This  again  is  interesting 
in  connection  with  a  certain  relation  between  dissociating 
power  and  another  property,  which  will  soon  be  discussed. 

The  dissociating  power  of  the  different  solvents  is  not 
simply  a  function  of  the  nature  of  the  solvents,  but  depends 
also  on  the  nature  of  the  dissolved  substance.  This  has 
long  been  known,  and  is  shown  very  clearly  by  work  on 
the  organic  acids  in  alcohol,  which  has  been  carried  out 

1  Zeit.  phys.  Chem.,  66,  179  (1906). 

4  Carnegie  Institution  of  Washington,  Publication  No.  230,  Chap.  II  (1915). 


NONAQUEOUS  AND  MIXED  SOLVENTS  209 

in  this  laboratory  during  the  past  two  years.  Wightman 
and  Wiesel1  found  that  the  conductivity  of  organic  acids 
hi  ethyl  alcohol  are  very  small  indeed;  being  hi  general 
several  hundred  times  less  than  hi  water  at  the  same  con- 
centration of  the  solutions.  This  was  confirmed  by  the 
work  of  Lloyd  and  Wiesel2  hi  which  the  conductivity  of 
about  forty  organic  acids  hi  ethyl  alcohol,  over  a  con- 
siderable range  of  dilution,  was  studied. 

Dissociating  Power  of  Solvents  and  Their  Dielectric 
Constant.  —  A  relation  of  importance  was  pointed  out 
first  by  Thomson,3  and  a  little  later  but  independently  by 
Nernst,4  between  the  dissociating  power  of  solvents  and 
their  dielectric  constants.  A  word  as  to  what  is  meant 
by  the  dielectric  constant  of  a  medium.  It  is  the  same 
property  of  media  that  was  called  by  Faraday  their 
"specific  inductive  capacity."  The  numerical  value  of  this 
constant  for  any  medium  determines  the  force  of  attrac- 
tion between  opposite  electrical  charges  separated  by  the 
medium  in  question.  Indeed,  the  dielectric  constant  has 
been  defined  as  the  relation  between  the  force  exerted 
between  two  charged  bodies  hi  a  vacuum  and  when  sepa- 
rated by  the  medium  hi  question.  This  can  be  seen  at 
once  if  we  consider  the  law  governing  such  phenomena  — 
the  law  of  Coulomb. 

Suppose  we  have  two  electrical  charges  which  we  will 
call  li  and  ^  at  a  distance  apart  r,  the  force  of  attraction 
or  repulsion,  /,  between  them  will  be  expressed  by  the 
equation: 

,1*1 

J"  r2  K 

in  which  K  is  the  dielectric  constant  of  the  medium  sepa- 
rating the  two  charges. 

We  see  at  once  that  K  being  hi  the  denominator,  the 

1  Journ.  Amer.  Chem.  Soc.,  36,  2243  (1914). 

2  Carnegie  Institution  of  Washington,  Publication  No.  230,  Chap.  VII 
(1915). 

1  Phil.  Mag.,  36,  320  (1893).  «  Zeit.  phys.  Chem.,  13,  531  (1894). 


210  THE  NATURE  OF  SOLUTION 

larger  its  value  the  smaller  the  electrostatic  force  acting 
between  the  charges,  —  the  force  is  inversely  propor- 
tional to  K. 

Why  should  there  be  any  relation  between  the  dielectric 
constants  of  media  and  their  own  power  to  dissociate 
molecules  dissolved  in  them?  The  author  has  been  best 
able  to  form  a  picture  of  this  relation  hi  the  following 
manner.  Take  for  example,  a  solvent  like  water,  which 
has  a  large  dielectric  constant.  The  force  of  attraction 
between  the  positively  charged  cation,  and  the  negatively 
charged  anion,  of  the  electrolyte  dissolved  in  the  water,  is 
inversely  proportional  to  the  dielectric  constant  of  the 
solvent.  This  being  large  in  the  case  of  water,  the  force 
of  attraction  between  the  positively  and  negatively 
charged  parts  is  small.  There  being  but  a  small  force 
to  hold  the  cation  and  the  anion  together  in  the  molecule, 
they  separate  or  undergo  electrolytic  dissociation  readily. 

If  the  dielectric  constant  of  the  medium  is  small,  as 
in  the,  case  of  the  hydrocarbons,  the  force  of  attraction 
between  the  oppositely  charged  parts  is  large,  and  they  do 
not  separate  or  undergo  electrolytic  dissociation  to  any 
appreciable  extent. 

Determination  of  the  Dielectric  Constants  of  Media.  — 
It  is  unsatisfactory  to  deal  with  any  quantity  and  not 
have  some  idea  how  it  is  measured.  A  number  of  methods 
have  been  proposed  and  used  for  determining  the  dielec- 
tric constants  of  liquids.  All  things  considered,  probably 
the  best  of  these  is  the  method  of  Drude.1  It  consists  in 
passing  electrical  waves  down  two  parallel  wires  surrounded 
by  the  medium  in  question,  and  measuring  the  length  of 
the  waves.  The  wave-length  is  inversely  proportional  to 
the  square  root  of  K,  or  the  dielectric  constant  of  the 
medium.  The  following  table  contains  the  dielectric  con- 
stants of  a  number  of  liquids,  many  of  which  are  good 
solvents,  and  also  then*  approximate  relative  dissociating 
powers. 

1  Zeit.  phys.  Chem.,  23,  267  (1897). 


NONAQUEOUS  AND  MIXED  SOLVENTS  211 

Dielectric  Dissociating 

Solvent  constant                                     power 

Hydrocyanic  acid  greater  than  water 

Hydrogen  dioxide  greater  than  water 

Formamide  84  greater  than  water 

Water  81.7  1 

Formic  acid  62.0  less  than  water 

Methyl  alcohol  35.3  about  $ 

Ethyl  alcohol  26.0  about  J 

Ammonia  22.0  about  | 

Acetone  20.7  about  £ 

Glycerol  16.5  $ 

Ethyl  ether  43.7  small 

Benzene  2.4  very  small 

Toluene  2.4  very  small 

That  the  relation  in  question  obtains,  at  least,  approxi- 
mately is  shown  by  the  above  results. 

J  Relation  Between  the  Dissociating  Power  of  Solvents 
and  Their  Own  Association.  —  A  relation  between  the 
dissociating  power  of  solvents  and  another  property  of  these 
liquids  was  pointed  out  by  Dutoit  and  Aston,1  which, 
while  it  does  not  hold  rigidly,  undoubtedly  contains  much 
of  value.  The  relation  hi  question  is  between  the  dis- 
sociating power  of  solvents  and  their  own  association;  and 
this  raises  the  question,  how  are  we  to  determine  the  degree 
of  association  of  any  substance  hi  the  liquid  state?  This 
has  been  done  by  several  independent  methods,  which  it 
would  lead  us  too  far  to  discuss  here  hi  any  detail.  One 
of  these  is  the  method  of  Ramsay  and  Shields,2  based  upon 
measurements  of  the  surface-tension  of  liquids,  and  which 
has  a  good  physical  foundation.  Another  is  the  purely 
empirical  method  of  Longinescu,3  which  consists  hi  deter- 
mining the  densities  and  boiling-points  of  liquids.  A  third 
method  has  been  suggested  by  Guye,4  which  is  based  upon 
the  measurement  of  the  refractivity  of  liquids  under  certain 
special  conditions.5 

The  significance  of  this  relation  can  be  seen  by  comparing 

Compt.  Rend.,  125,  240  (1897). 
Zeit.  phys.  Chem.,  12,  433  (1893). 
Journ.  Chim.  Phys.,  1,  289  (1903). 
Ann.  Chim.  Phys.  [6],  21,  211  (1890). 

For  details  see  Elements  of  Physical  Chemistry,  by  the  author,  4th  edi- 
tion, p.  145.    (The  Macmillan  Co.) 


212  THE  NATURE  OF  SOLUTION 

the  association  factors  for  a  number  of  liquids,  at  a  given 
temperature,  with  the  dissociating  powers  of  these  liquids 
acting  as  solvents. 

Association  Dissociating 

Solvent  factor  power 

Formamide  6.2  greater  than  water 

Water  3.8  1 

Methyl  alcohol  3.4  about  < 

Ethyl  alcohol  2.7  about  i 

Glycerol  1.8  about  j- 

Acetone  1.3  about  f 

In  making  this  comparison  we  are  limited  to  those  liquids 
of  which  the  association  factors  have  been  ascertained. 

This  relation  between  the  dissociating  power  of  sol- 
vents and  their  own  association  is  not  a  quantitative  one, 
as  has  already  been  pointed  out;  but  it  is  rather  more 
than  qualitative.  Other  things  being  equal,  those  solvents 
which  are  the  most  associated  have  the  greatest  dissocia- 
ting power.  It  might  also  be  pointed  out  in  passing  that 
many  of  the  best  solvents  are  strongly  associated  com- 
pounds. 

Effect  of  Rise  in  Temperature  on  the  Association  of 
Liquids  and  on  Their  Dissociating  Power.  —  When  we  were 
studying  the  dissociation  of  electrolytes  as  measured  by 
conductivity,  it  was  pointed  out  that  rise  in  temperature 
diminishes  the  dissociation  of  a  solution  of  an  electrolyte. 
This  raises  the  question,  what  is  the  effect  of  rise  in  tem- 
perature on  the  association  of  liquids? 

It  is  well  known  that  rise  in  temperature  tends  to  break 
down  complexes  hi  general;  and  this  has  been  shown  to 
be  the  case  with  associated  molecules.  The  association 
of  liquids  over  a  wide  range  in  temperature,  extending 
from  ordinary  temperatures  almost  to  the  critical  tem- 
perature of  the  liquid  in  question,  was  measured  by  Ramsay 
and  Shields,  with  the  result  that  the  higher  the  temperature 
the  less  the  liquid  is  associated.  This  is,  then,  another  re- 
lation between  the  association  of  liquids  and  their  power 
to  break  down  molecules  into  ions. 

Results  in  Mixed  Solvents.  —  A  fairly  large  amount  of 


NONAQUEOUS  AND  MIXED  SOLVENTS  213 

work  has  been  done,  especially  in  this  laboratory,  on  the 
condition  of  substances  in  mixtures  of  solvents. 

The  observation  was  made  by  Zelinsky  and  Krapiwin,1 
that  certain  salts  when  dissolved  hi  a  certain  mixture  of 
methyl  alcohol  and  water,  showed  lower  conductivities  than 
when  dissolved  in  pure  methyl  alcohol]  notwithstanding  the 
fact  that  the  same  salts,  under  the  same  condition  of  dilu- 
tion, showed  much  lower  conductivity  hi  methyl  alcohol 
than  hi  water.  This  was  indeed  a  remarkable  fact,  and 
merited  careful  and  extensive  investigation.  The  follow- 
ing discussion  is  taken  from  Publication  of  the  Carnegie 
Institution  of  Washington,  No.  210,  pp.  162-170. 2 

Mixtures  of  Water  and  the  Alcohols.  —  The  study  of 
the  conductivities  and  dissociations  hi  pure  solvents  was 
extended  hi  this  laboratory  to  mixed  solvents,  and  the 
results  have  been  published  hi  monographs  Nos.  80  and 
180  of  the  Carnegie  Institution  of  Washington. 

The  first  investigation  was  carried  out  by  Lindsay 3 
who  worked  hi  water,  hi  methyl,  ethyl,  and  propyl  alcohols, 
and  hi  mixtures  of  these  solvents  with  one  another.  He 
found  that  hi  certain  mixtures  of  the  alcohols  with  water, 
the  conductivity  of  the  dissolved  salt  was  less  than  hi  the 
pure  alcohol.  The  conductivity  curves  in  mixtures  of 
methyl  alcohol  and  water  showed  very  distinct  minima, 
and  a  conductivity  minimum  was  also  frequently  found 
in  mixtures  of  ethyl  alcohol  and  water. 

A  possible  explanation  of  the  results  hi  mixtures  of  the 
alcohols  with  water  is  that  each  solvent  diminishes  the 
association  of  the  other.  Since  the  dissociating  power  of  a 
solvent  is  hi  general  greater  the  larger  its  own  association, 
it  follows  that  whatever  would  decrease  the  association  of  a 
liquid  would  decrease  its  power  to  dissociate  electrolytes 
dissolved  in  it.  The  question  is,  does  one  associated  liquid 
diminish  the  association  of  another  associated  liquid? 

1  Zeit.  phys.  Chem.,  21,  35  (1896). 

8  See  also  Journal  of  the  Franklin  Institute,  Nov.  and  Dec..  1913. 

1  Amer.  Chem.  Journ.,  28,  329  (1902). 


214  THE  NATURE  OF  SOLUTION 

An  associated  liquid  tears  down  the  molecules  of  an 
electrolyte  dissolved  in  it  into  simpler  parts  or  ions;  and 
it  might  be  expected  that  such  a  liquid  would  tear  down 
the  molecules  of  another  associated  liquid,  a  non-electrolyte, 
not  into  charged  parts  or  ions,  but  into  simpler  molecules. 
Alcohol  and  water  are  associated  liquids,  as  has  been 
shown  by  the  surface-tension  measurements  of  Ramsay  and 
Shields.1  Do  these  diminish  the  association  of  one  another? 

That  this  is  the  case  was  shown  by  Murray.2  He 
worked  with  the  associated  liquids,  water,  formic  acid, 
and  acetic  acid.  He  determined  the  molecular  weight  of 
each  of  these  liquids  in  the  other  two,  and  found  that  their 
molecular  weights  became  smaller  the  more  dilute  the 
solutions.  This  showed  that  the  solvent,  i.e.,  the  liquid 
present  hi  the  larger  quantity,  was  tearing  down  the 
molecular  complexes  of  the  dissolved  liquid  or  the  one 
present  in  smaller  quantity. 

That  there  is  a  diminution  in  the  association  of  one  asso- 
ciated liquid  by  another  associated  liquid  was  shown  for  the 
above-named  substances  and  made  highly  probable  for  others. 

That  this  was  not  the  entire  explanation  of  the  nature 
of  the  conductivity  curves  in  mixtures  of  certain  alcohols 
with  water,  was  brought  out  by  the  next  investigation  in 
this  field,  carried  out  by  Carroll.3  He  compared  the 
conductivity  curves  of  electrolytes  dissolved  in  these  sol- 
vents with  the  fluidity  curves  of  the  mixtures  of  the 
two  liquids  in  question,  and  found  that  the  two  sets  of 
curves  were  very  similar.  The  minima  in  the  two  cases 
occurred  in  the  same  mixture  of  the  two  liquids.  A  care- 
ful comparison  of  the  two  sets  of  phenomena  led  us  to 
conclude  that  the  conductivity  maxima  are  largely  due  to 
the  decrease  in  fluidity  which  takes  place  on  mixing  the 
two  solvents.  The  diminished  fluidity,  or  increased  vis- 
cosity, would  cause  the  ions  to  move  more  slowly,  and 
hence  decrease  the  conductivity. 

1  Zeit.  phys.  Ckem.,  12,  433  (1893). 

8  Ajner.  Chem.  Jour.,  30, 193  (1903).  «  Ibid.,  32,  521  (1904). 


NONAQUEOUS  AND  MIXED  SOLVENTS  215 

At  the  end  of  the  work  done  by  Carroll,  it  seemed 
justifiable  to  conclude  that  the  conductivities  of  binary 
electrolytes  in  such  solvents  as  those  already  considered, 
are  inversely  proportional  to  the  coefficients  of  viscosity 
of  the  solvent  and  are  directly  proportional  to  the  associa- 
tion of  the  solvent.  Bassett1  showed  that  silver  nitrate  hi 
mixtures  of  methyl  alcohol  and  water  gave  a  conductivity 
minimum  at  both  0°  and  25°;  also  that  the  effect  of  one 
solvent  on  the  other  was  greater  at  0°  than  at  25°.  This 
would  be  expected,  since  the  dissociation  diminishes  with 
rise  in  temperature,  and  each  solvent  would  probably 
diminish  the  association  of  the  other  less,  the  smaller  its 
own  association  or  the  higher  its  temperature. 

Mixtures  of  Acetone  with  Water  and  the  Alcohols.  — 
Bingham2  measured  not  only  the  conductivities,  but  also 
the  viscosities  of  a  number  of  solvents  and  solutions  La 
these  solvents.  He  found  minima  hi  the  conductivity 
curves  hi  mixtures  of  acetone  and  water.  The  distinctly 
new  feature  brought  out  by  the  work  of  Bingham  was  that 
lithium  and  calcium  nitrates  hi  mixtures  of  acetone  with 
methyl  and  ethyl  alcohols  showed  a  decided  maximum 
in  the  conductivity  curves.  This  must  be  due  either  to  an 
increase  in  dissociation  in  such  mixtures,  increasing  the 
number  of  ions  present,  and  consequently  increasing 
the  conductivity,  or  it  must  be  due  to  a  diminution  in  the 
complexity  of  the  solvates  around  the  ions,  increasing  their 
velocities.  The  dissociation  was  measured  in  the  mixtures 
in  question  and  found  not  to  account  for  the  phenomenon. 
This  eliminates  increase  hi  dissociation  and  leaves  the 
other  alternative,  diminution  hi  the  complexity  of  the  sol- 
vate,  to  account  for  the  phenomenon. 

The  ion  must  take  with  it  through  the  solvent  any  mole- 
cules of  the  liquid  with  which  it  might  be  combined.  This 
would  increase  the  effective  mass  and  diminish  its  veloc- 
ity. Anything  which  would  diminish  the  complexity  of  the 
solvate  about  the  ion  would  increase  its  velocity,  and  con- 

1  Amer.  Chem.  Journ.,  32,  409  (1904).  «  Ibid.,  34,  481  (1905). 


216  THE  NATURE  OF  SOLUTION 

sequently  the  conductivity.  We  must  therefore  conclude 
that  the  solvates  in  those  mixtures  of  acetone  with  the 
alcohols  are  simplest  where  the  conductivity  is  the  greatest. 

Rouiller1  studied  both  the  velocities  of  the  ions  and  the 
conductivities  of  electrolytes  in  mixtures  of  acetone  with 
the  alcohols.  Silver  nitrate  in  methyl  alcohol  and  acetone 
gave  a  decided  maximum  of  conductivity.  His  work  on 
the  velocities  of  the  ions  in  these  mixtures  indicated  that 
the  above  explanation  of  the  maxima  offered  by  Jones  and 
Bingham  was  correct;  there  is  a  change  hi  the  complexity 
of  the  solvate  about  the  ion. 

McMaster2  extended  the  work  hi  the  same  solvents 
—  water,  methyl  alcohol,  ethyl  alcohol,  and  acetone — 
and  hi  mixtures  of  these  with  each  other.  He  found 
conductivity  results  of  the  same  general  character  as 
those  obtained  by  the  earlier  workers.  Conductivity 
minima  were  found  hi  mixtures  of  the  alcohols  with  water 
and  acetone  with  water.  Conductivity  maxima  were 
obtained  with  lithium  bromide  in  mixtures  of  methyl  or 
ethyl  alcohol  with  acetone.  Cobalt  chloride  in  mixtures 
of  acetone  with  ethyl  alcohol  also  showed  a  maximum. 
Jones  and  McMaster  reached  the  same  conclusion  from 
their  work  that  had  been  reached  by  Jones  and  Bingham. 
Change  in  the  complexity  of  the  solvate  formed  by  the  ion 
in  different  mixtures  of  solvents  is  an  important  factor 
in  determining  the  conductivity  maxima. 

A  point  of  interest  brought  out  by  the  work  of  McMas- 
ter is  in  connection  with  the  temperature  coefficients  of 
conductivity  in  non-aqueous  solutions.  The  bearing  of  tem- 
perature coefficients  of  conductivity  on  the  solvate  theory 
of  solution  has  already  been  discussed.  With  rise  in  tem- 
perature the  hydrates  about  the  ions  become  simpler. 
The  mass  and  probably  the  size  of  the  ion  thus  becomes  less, 
and  it  moves  faster  the  higher  the  temperature,  thus 
increasing  the  conductivity. 

McMaster  found  that  cobalt  chloride  in  certain  mixtures 

»  Amer.  Chem.  Journ.,  36,  427  (1906).  8  76id.,  36,  325  (1906). 


NONAQUEOUS  AND  MIXED  SOLVENTS  217 

of  acetone  with  the  alcohols  showed,  at  ordinary  tempera- 
tures, negative  temperature  coefficients  of  conductivity. 
What  does  this  mean?  The  solvent  becomes  less  viscous 
with  rise  hi  temperature,  thus  increasing  the  velocity  of 
the  ions;  and  the  solvates  become  simpler,  which  also 
increases  the  velocity  with  which  the  ions  move. 

With  rise  in  temperature,  on  the  other  hand,  the 
association  of  the  solvent,  and  consequently  its  dissocia- 
ting power,  becomes  less. 

The  above  two  influences  work  counter  to  one  another. 
Negative  temperature  coefficients  of  conductivity  mean 
that  the  latter  influence  predominates.  The  alcohols 
used  and  acetone  are  highly  associated  liquids.  Rise  hi 
temperature  diminishes  their  association  and  consequently 
their  dissociating  power. 

A  solution  of  cobalt  chloride  in  a  75  per  cent  mixture  of 
acetone  with  methyl  alcohol,  the  solution  being  g-J^  nor- 
mal, had  a  zero  temperature  coefficient  of  conductivity. 

Viscosity  and  Atomic  Volume.  —  A  number  of  points  of 
interest  were  brought  out  by  the  next  investigator,  Veazey.1 
He  worked  with  solutions  of  salts  in  water,  methyl  alcohol, 
ethyl  alcohol,  acetone,  and  in  binary  mixtures  of  these  liquids 
with  one  another.  The  minimum  hi  conductivity  was  found 
to  be  a  more  general  phenomenon  than  had  been  supposed 
from  the  earlier  work.  It  had  long  been  known  that  mix- 
tures of  methyl  alcohol  and  water  or  ethyl  alcohol  and  water, 
are  more  viscous  than  either  of  the  pure  solvents  alone.  A 
rational  explanation  of  this  phenomenon  was  suggested  — 
alcohol  and  water  are  strongly  associated  liquids.  When 
two  associated  liquids  are  mixed  each  diminishes  the  asso- 
ciation of  the  other.  The  larger  molecules  are  thus  broken 
down  into  smaller  molecules,  which  increases  the  frictional 
surfaces  when  these  molecules  move  over  one  another  as  they 
do  hi  viscous  flow.  The  result  would  be  to  increase  the 
viscosity  of  the  mixture  over  that  of  either  pure  solvent. 

*  Amer.  Chem.  Jaurn.,  37,  405  (1907).  Zeit.  phya.  Chem.,  61,  641  (1908); 
62,  44  (1908). 


218  THE  NATURE  OF  SOLUTION 

Maxima  in  the  conductivity  of  electrolytes  in  the  mixed 
solvents  were  shown  to  correspond  to  maxima  in  the 
fluidity  of  the  mixed  solvents.  Maxima  in  fluidity  are 
probably  due  to  an  increase  in  the  size  of  the  molecules 
of  the  solvent,  due  to  a  combination  of  one  liquid  with 
the  other.  This  would  diminish  the  viscosity  and  conse- 
quently increase  the  velocity  of  the  ions,  which  would 
increase  the  conductivity.  This  factor  must  also  be  taken 
into  account  in  explaining  conductivity  maxima. 

The  temperature  coefficients  of  conductivity  in  the 
above-named  mixtures  of  liquids  with  water  are  a  maxi- 
mum in  the  25  and  50  per  cent  mixtures.  These  are  just 
about  the  mixtures  in  which  the  solvents  have  the  least 
association.  The  molecules  of  the  solvent  being  in  the 
simplest  condition  would  be  most  favorable  for  chemical 
action.  In  such  mixtures  the  solvents  probably  combine 
to  the  greatest  extent  with  the  dissolved  substance  —  the 
solvation  is  at  a  maximum.  The  effect  of  rise  in  tem- 
perature breaking  down  these  solvates  would  therefore  be 
a  maximum  where  solvation  is  a  maximum.  Solutions 
of  potassium  sulphocyanate  have  greater  conductivity  in 
acetone  than  in  water.  This  was  shown  to  be  due  to  the 
greater  fluidity  of  the  acetone. 

This  same  salt  when  dissolved  hi  water  lowers  the  vis- 
cosity of  the  water.  Certain  salts  of  potassium  and  salts 
of  rubidium  and  caesium  were  practically  the  only  ones 
known  at  that  to  lower  the  viscosity  of  water.  In  the 
case  of  certain  salts  of  potassium  the  positive  effect  of  the 
anion  on  the  viscosity  of  water  may  more  than  offset  the 
negative  effect  of  the  potassium  ion. 

The  following  explanation  of  the  above-named  phenom- 
enon was  suggested.  If  the  atomic  volume  of  the  ions 
dissolved  in  the  solvent  were  larger  than  the  molecular 
volume  of  the  solvent,  the  larger  ions  would  diminish  the 
size  of  the  frictional  surfaces  coming  in  contact  and  would 
lower  the  viscosity. 

It  is  well  known  that  potassium,  rubidium,  and  caesium 


NONAQUEOUS  AND  MIXED  SOLVENTS  219 

occupy  the  maxima  on  the  atomic-volume  curve,  and  have 
much  larger  atomic  volumes  than  any  other  known  ele- 
ments. Potassium  has  a  smaller  atomic  volume  than 
rubidium,  and  rubidium  than  caesium.  Potassium  chlo- 
ride lowers  the  viscosity  of  water  less  than  rubidium 
chloride,  and  rubidium  chloride  less  than  caesium  chloride. 

If  we  study  the  salts  which  raise  the  viscosity  of  water, 
we  will  find,  in  general,  that  the  amount  of  increase  hi 
the  viscosity  bears  a  relation  to  the  atomic  or  ionic  vol- 
umes of  the  dissolved  substances.  Smaller  ions  tend  to 
increase  the  viscosity  of  water  more  than  larger  ones.  It 
would  therefore  seem  that  the  above  explanation  contains 
a  large  element  of  truth. 

Dissociation  in  Nonaqueous  Solvents.  —  The  problem 
of  measuring  dissociation  in  nonaqueous  solvents  is  a 
difficult  one.  The  freezing-point  method  is  frequently 
not  applicable.  Many  common  solvents,  such  as  the 
alcohols,  freeze  at  temperatures  which  are  too  widely  re- 
moved from  the  ordinary  temperature  of  the  laboratory  to 
be  measured  with  sufficient  accuracy.  The  boiling-point 
method  could  be  used  only  with  fairly  concentrated  solu- 
tions. Dilute  solutions  produce  such  a  slight  rise  in  the 
boiling-point  that  this  small  quantity  cannot  be  measured 
with  a  very  high  degree  of  accuracy.  The  boiling-point 
method  has  the  further  disadvantage  of  being  so  largely 
affected  by  slight  changes  in  the  barometer. 

The  hope  of  measuring  dissociation  in  nonaqueous 
solvents  in  general  seemed  to  rest  hi  the  conductivity 
method.  This  method  as  ordinarily  applied  would  not  be 
satisfactory.  The  dilution  at  which  complete  dissociation 
would  be  reached  in  such  solvents  is  so  great  that  the 
Kohlrausch  method  in  any  such  form  as  he  left  it  could  not 
be  applied  to  the  problem. 

The  conductivity  method  was  greatly  improved  by 
Kreider1;  the  greatest  improvement  being  in  the  form  of 
cell  employed.  With  the  improved  method  Kreider  studied 

1  Amer.  Chem.  Journ.,  45,  282  (1911). 


220  THE  NATURE  OF  SOLUTION 

the  dissociations  of  a  number  of  salts  in  methyl  and  ethyl 
alcohols  and  in  mixtures  of  these  solvents  with  water.  He 
measured  the  conductivities  of  solutions  as  dilute  as  100,- 
000  liters  and  found  the  following  relation  to  hold: 

fj,m  methyl  alcohol 


ethyl  alcohol 


=  constant. 


When  a  salt  is  equally  dissociated  by  each  of  two  sol- 
vents, for  the  same  concentration  of  the  salt  there  is  the 
same  number  of  ions  in  the  two  solutions.  Conductivity  is 
a  function  of  the  number  of  the  ions  and  their  velocities. 
When  the  number  of  the  ions  is  constant,  as  in  this  case, 
conductivity  is  a  function  of  the  relative  velocities  of  the 
ions.  The  velocity  of  an  ion  is  conditioned  by  its  mass 
and  volume  and  by  the  fluidity  of  the  solvent.  If  the 
mass  and  volume  of  the  ions  in  the  two  solvents  are  con- 
stant, the  velocities  of  the  ions  should  vary  as  the  fluidi- 
ties of  the  solvents.  The  ratio  between  the  values  of  /*„ 
in  the  two  solvents  should  be  the  same  as  the  ratio  between 
the  fluidities  of  these  solvents.  This  was,  however,  found 
not  to  be  the  case.  The  bearing  of  this  fact  on  the  condi- 
tion of  the  ions  in  the  two  solvents  in  question  is  impor- 
tant. This  shows  that  the  mass  and  probably  the  volume 
of  the  solvated  ion  must  differ  in  the  two  solvents. 

The  ratio  between  the  values  of  /*«>  for  a  salt  in  the  two 
solvents,  compared  with  the  ratio  between  the  fluidities  of 
the  two  solvents,  would  give  an  approximate  idea  of  the  rela- 
tive solvation  of  the  ions  in  the  two  solvents  in  question. 

This  method  will  be  still  further  applied  to  the  problem 
of  solvation  in  non-aqueous  solvents. 

Ternary  Mixtures  of  the  Alcohols  with  Water,  —  Mahin1 
studied  electrolytes  in  ternary  mixtures  of  the  alcohols 
with  water,  and  obtained  results  of  the  same  general  charac- 
ter as  those  found  in  binary  mixtures  of  these  solvents. 
He  then  took  up  work  in  binary  mixtures,  one  constituent 
being  acetone.  Acetone  was  studied  primarily  because  in 

1  Amer.  Chem.  Journ.,  41,  433  (1909);  Zeit.  phys.  Chem.,  69,  389  (1909). 


NONAQUEOUS  AND  MIXED  SOLVENTS  221 

many  of  its  properties  it  is  an  exceptional  solvent.  Sub- 
stances dissolved  in  acetone  are  largely  polymerized,  and 
acetone  has  at  the  same  time  considerable  dissociating 
power.  Furthermore,  acetone  is  a  solvent  with  small  vis- 
cosity, and  it  was  desired  to  see  whether  the  relations  found 
for  solvents  with  larger  viscosity  would  hold  here.  The 
curve  for  conductivity  and  for  fluidity  were  worked  out 
and  the  two  compared. 

It  was  found  that  the  product  of  molecular  conduc- 
tivity and  viscosity  is  nearly  a  constant  at  complete  dis- 
sociation. This  means  that  for  completely  dissociated 
solutions  in  acetone  the  curves  of  molecular  conductivity 
are  similar  to  those  of  fluidity  —  conductivity  being  in- 
versely proportional  to  viscosity.  This  relation  is  of 
interest  in  that  it  holds  hi  a  solvent  with  such  small  vis- 
cosity as  acetone. 

Relations  such  as  those  referred  to  above  having  been 
found  to  hold  hi  a  solvent  with  such  small  viscosity  as 
acetone,  the  question  arose,  do  such  relations  obtain  in  a 
highly  viscous  solvent  like  glycerol?  Glycerol  not  only 
has  a  very  high  viscosity,  but  is  an  excellent  solvent,  and 
has  a  large  dielectric  constant,  which  means  that  it  has 
considerable  dissociating  power.  Glycerol  is  somewhat 
strongly  associated,  which  also  indicates  considerable  dis- 
sociating power. 

Glycerol  and  Mixtures  with  the  Alcohols  and  Water.  — 
The  first  investigation  hi  glycerol  as  a  solvent  was  carried 
out  by  Schmidt.1  He  measured  the  conductivities  of  solu- 
tions of  certain  salts  hi  glycerol,  and  hi  mixtures  of  glycerol 
with  water  and  with  methyl  and  ethyl  alcohols.  The  con- 
ductivities were  measured  at  different  temperatures.  The 
most  striking  relation  noted  was  the  enormous  magnitude 
of  the  temperature  coefficients  of  conductivity  of  electro- 
lytes dissolved  in  glycerol.  This  was  shown  to  be  due  to 
the  rapid  decrease  hi  the  viscosity  of  glycerol  with  rise  in 
temperature. 

1  Amer.  Chem.  Journ.,  42,  37  (1909). 


222  THE  NATURE  OF  SOLUTION 

It  was  shown  that  when  glycerol  is  mixed  with  water 
or  the  alcohols,  there  is  a  breaking  down  of  the  association 
of  each  solvent  by  the  other,  and  a  consequent  diminution 
hi  the  dissociating  power.  Solutions  of  potassium  iodide 
in  25  and  50  per  cent  mixtures  of  glycerol  and  water  were 
less  viscous  than  the  solvents  themselves.  This  salt  does 
not  lower  the  viscosity  of  glycerol,  but  of  the  mixtures. 
The  meaning  of  negative  viscosity  effects  was  discussed 
in  the  work  of  Veazey.  While  Schmidt  did  not  study  any 
salt  which  lowers  the  viscosity  of  pure  glycerol,  he  found 
that  the  effect  of  the  salt  on  the  viscosity  of  pure  glycerol 
was  inversely  as  the  molecular  volume  or  atomic  volumes 
of  the  constituents  of  the  salt.  This  was  in  keeping  with 
the  explanation  offered  by  Jones  and  Veazey  to  account 
for  the  changes  in  the  viscosity  of  the  solvent  by  the  dis- 
solved substance.  A  comparison  of  the  conductivity  and 
fluidity  curves  shows  that  the  two  run  nearly  parallel. 
Although  glycerol  has  about  1,000  times  the  viscosity  of 
methyl  alcohol,  yet,  from  the  work  of  Schmidt,  the  same 
general  relations  obtain  here  that  hold  for  the  far  less 
viscous  solvents. 

The  work  of  Schmidt  was  continued  by  Guy.1  He 
worked  with  a  much  larger  number  of  salts,  and  over  the 
temperature  range  25°  to  75°.  He  studied  not  only  solu- 
tions in  glycerol,  but  in  mixtures  of  glycerol  with  water, 
with  methyl,  and  with  ethyl  alcohols. 

Guy  found  also  enormous  temperature  coefficients  of 
conductivity.  This  may  be  due  to  either  of  two  causes: 
a  change  hi  dissociation  with  rise  hi  temperature,  or  a 
change  in  the  velocity  of  the  ions.  We  know  the  order 
of  magnitude  of  the  change  in  dissociation  with  rise  hi  tem- 
perature, and  it  is  small.  The  chief  cause  of  the  large 
temperature  coefficients  of  conductivity  in  glycerol  is, 
then,  an  increase  in  the  velocities  with  which  the  ions 
move.  As  we  have  seen,  this  may  be  due  to  a  decrease 
in  the  viscosity  of  the  solvent  with  rise  in  temperature, 

1  Amer.  Chem.  Journ.,  46,  131  (1911). 


NONAQUEOUS  AND  MIXED  SOLVENTS  223 

or  may  be  caused  by  a  breaking  down  of  complex  solvates 
about  the  ions. 

While  the  viscosity  of  glycerbl  decreases  rapidly  with 
rise  in  temperature,  this  alone  would  not  account  for  the 
magnitude  of  the  temperature  coefficients  of  conductivity 
of  glycerol  solutions.  There  seems  to  be  good  evidence 
for  the  formation  of  glycerolates  hi  solutions  hi  glycerol. 
The  temperature  coefficients  of  conductivity  hi  glycerol 
are  greater  at  high  than  at  low  dilution.  Jones  has  pointed 
out  that  this  would  be  expected  from  the  solvate  theory. 
The  more  dilute  the  solution  the  more  complex  the  solvate; 
the  more  complex  the  solvate  the  greater  the  change  hi  its 
complexity  with  rise  hi  temperature. 

Further,  salts  of  calcium,  strontium,  and  barium  have 
larger  temperature  coefficients  of  conductivity  than  those 
of  sodium,  potassium,  and  ammonium.  The  former  are 
strongly  hydrated,  the  latter  weakly  hydrated  substances. 
It  would  seem  that  the  former  are  more  strongly  glycerol- 
ated  than  the  latter.  Salts  which  have  approximately 
the  same  hydrating  power  have  temperature  coefficients 
of  conductivity  hi  glycerol  of  the  same  order  of  magnitude, 
indicating  the  same  order  of  magnitude  of  glycerolation. 
Work  in  the  mixed  solvents  indicates  that  water  diminishes 
the  association  of  glycerol. 

Solutions  of  salts  in  glycerol  have  in  general  greater 
viscosity  than  pure  glycerol.  Guy,  however,  found  marked 
exceptions  to  this  relation.  Salts  of  rubidium  lowered  the 
viscosity  of  glycerol.  Ammonium  bromide  and  iodide  also 
lowered  the  viscosity  of  this  solvent.  That  rubidium 
should  lower  the  viscosity  of  glycerol  is  hi  keeping  with 
what  was  found  in  aqueous  solutions.  Salts  of  rubidium 
and  caesium  and  some  salts  of  potassium  lowered  the  vis- 
cosity of  water.  This  has  already  been  explained  as  due 
to  the  large  atomic  volumes  of  these  elements.  The  same 
explanation  holds  for  solutions  hi  glycerol. 

Davis1  continued  the  work  of  Guy,  studying  especially 

1  Zeit.  phys.  Chem.,  81,  68  (1912). 


224  THE  NATURE  OF  SOLUTION 

the  effect  of  salts  on  the  viscosity  of  glycerol.  He  repeated 
the  work  with  ammonium  iodide  and  obtained  the  same 
result  that  had  been  earlier  found  by  Guy.  He  studied 
rubidium  chloride,  bromide,  iodide,  and  nitrate,  and 
showed  that  these  lowered  the  viscosity  of  glycerol.  The 
rubidium  salts  lower  the  viscosity  of  glycerol  to  such  an 
extent  that  they  appreciably  increase  their  own  conduc- 
tivity in  this  solvent. 

Comparing  the  effects  of  the  chloride,  bromide,  and 
iodide  of  rubidium  on  the  viscosity  of  glycerol,  Davis 
found  that  the  chloride  has  the  least  effect,  the  bromide 
next,  the  iodide  the  greatest.  He  showed  that  this  was  in 
the  same  order  as  the  molecular  volumes  of  the  salts  in 
question.  The  results  obtained  with  glycerol  were,  then, 
analogous  to  those  obtained  with  water,  both  with  respect 
to  viscosity  and  solution. 

Rubidium  Salts  in  Mixtures  of  Acetone  and  Water.  — 
Davis  and  Hughes1  studied  the  conductivities  and  vis- 
cosities of  solutions  of  rubidium  salts  in  mixtures  of  ace- 
tone and  water,  from  the  standpoint  of  the  theory  of 
viscosity  proposed  by  Jones  and  Veazey.  They  found  that 
the  viscosities  of  the  mixture  were  lowered  when  it  did  not 
contain  an  excess  of  acetone;  the  effect,  however,  being 
less  than  in  water  or  in  glycerol. 

Work  in  Fonnamide.  —  Davis  and  Putnam2  studied  the 
viscosities  of  solutions  in  formamide  as  a  solvent.  This 
was  of  interest  in  that  formamide  has  a  higher  dielectric 
constant  than  water,  and  would  therefore  be  expected 
to  have  greater  dissociating  power.  Further,  it  is  a  very 
strongly  associated  solvent,  being  more  associated  than 
water,  and  this  also  would  indicate  greater  dissociating 
power. 

A  specially  devised  apparatus  for  distilling  under  di- 
minished pressure  enabled  the  investigators  to  prepare 
formamide  with  a  conductivity  comparable  with  that  of 

1  Zeit.  phys.  Chem.,  85,  513  (1913). 

a  Carnegie  Institution  of  Washington,  Publication  No.  230,  Chap.  II  (1915). 


NONAQUEOUS  AND  MIXED  SOLVENTS  225 

conductivity  water.  They  showed  that  this  solvent  has 
greater  dissociating  power  than  water,  and  that  those  salts 
which  form  hydrates  with  water  are  solvated  also  in  for- 
mamide. 

The  temperature  coefficients  of  conductivity  hi  for- 
mamide  are  what  would  be  expected  from  the  coefficients 
of  hydrated  salts  in  aqueous  solutions. 

Viscosities  of  Caesium  Salts.  —  Having  obtained  a 
liberal  supply  of  a  caesium  salt  through  the  co-operation 
of  Professor  Howe  of  Washington  and  Lee  University, 
Davis 1  was  able  to  study  the  effect  of  caesium  salts  on  the 
viscosity  of  water  and  other  solvents.  Caesium,  as  is 
well  known,  occupies  the  highest  maximum  on  the  atomic 
volume  curve,  having  the  largest  atomic  volume  of  all  the 
elements.  In  terms  of  the  theory  proposed  by  Jones  and 
Veazey,  caesium  salts  should  lower  the  viscosity  of  sol- 
vents even  more  than  salts  of  rubidium,  and  such  is  the 
fact. 

1  Carnegie  Institution  of  Washington,  Publication  No.  230,  Chap.  I  (1915). 


CHAPTER  XII 

COLLOIDAL   SOLUTIONS 

Historical  Sketch.  —  There  are  few  phases  of  solution 
which  have  recently  attracted  so  much  attention  as  col- 
loids. We  hear  continually  of  colloidal  solutions  and  of 
the  remarkable  properties  possessed  by  matter  in  the  col- 
loidal state. 

While  this  subject  has  come  to  the  front  only  in  the 
last  few  years,  matter  hi  the  colloidal  state  has  been 
known  for  over  half  a  century.  The  Italian  chemist 
Selmi,  as  early  as  1844,  recognized  the  difference  between 
true  solutions  and  apparent  solutions  of  such  substances 
as  sulphur,  Berlin  blue,  and  the  like.  He  called  the  latter 
pseudo-solutions,  and  recognized  certain  fundamental  dif- 
ferences between  these  and  true  solutions.  He  pointed 
out  that  when  pseudo-solutions  are  formed  no  change  in 
temperature  is  produced.  Further,  there  is  neither  con- 
traction nor  expansion  of  the  liquid,  and  when  salts  are 
added,  the  colloid  is  precipitated.  He  observed  that  in 
such  precipitations,  a  part  of  the  precipitant  is  carried 
down  with  the  colloid. 

Selmi  assumed  that  in  pseudo-solutions  the  substance 
was  in.  the  state  of  an  emulsion  or  a  suspension. 

Work  of  Graham.  —  The  first  to  study  colloids  exten- 
sively was  the  English  chemist,  Graham.  He  found  that 
those  compounds  which  readily  form  crystals,  diffuse 
rapidly  through  membranes  made  of  vegetable  parchment; 
while  non-crystallizable  or  amorphous  substances  either  do 
not  diffuse  at  all  through  such  membranes,  or  diffuse 
very  slowly  through  them.  The  former  Graham  called 
crystalloids,  and  the  latter,  colloids.  He  used  diffusion  for 


COLLOIDAL  SOLUTIONS  227 

separating  the  one  class  from  the  other  —  crystalloids 
from  colloids  —  and  likened  these  separations  by  diffusion 
to  the  separations  based  upon  the  different  volatility  of 
substances. 

Graham  saw  the  importance  of  colloids  not  only  for 
inanimate,  but  for  living  matter.  This  can  be  seen  best 
by  quoting  his  own  words.1  "The  colloidal  is,  hi  fact,  a 
dynamical  state  of  matter;  the  crystalloidal  being  the 
statical  condition.  The  colloid  possesses  energia.  It  may 
be  looked  upon  as  the  probable  primary  source  of  the 
force  appearing  in  the  phenomena  of  vitality.  To  the 
gradual  manner  hi  \vhich  colloidal  changes  take  place  (for 
they  always  demand  tune  as  an  element),  may  the  char- 
acteristic protraction  of  chemical-organic  changes  also  be 
referred." 

This  is  prophetic  of  what  is  now  supposed  to  be  the 
importance  of  colloids  for  the  life  process,  and  for  the 
normal  functions  of  living  matter. 

Diffusion  Experiments  of  Graham.  —  Graham  carried 
out  his  diffusion  experiments  with  comparatively  crude 
apparatus.  The  bottom  of  a  wide  glass  cylinder  or  dish 
was  covered  with  vegetable  parchment,  and  this  was 
floated  on  water.  The  substance  was  placed  hi  the  vessel 
and  separated  from  the  water  by  the  membrane.  The 
crystalloids  passed  through  the  membrane,  and  the  col- 
loids either  did  not  pass  through  at  all,  or  passed  through 
only  very  slowly. 

A  few  of  the  results  obtained  by  Graham  will  give  an 
idea  of  the  relative  rates  at  which  crystalloids  and  col- 
loids diffuse.  Representing  as  unity  the  amount  of  sodium 
chloride  which  diffuses  through  the  parchment  paper  hi 
twenty-four  hours  at  10°  to  15°,  Graham2  found  that  the 
relative  rates  at  which  other  substances  will  diffuse  are 
those  given  in  the  following  table. 

1  PhU.  Trans.,  151,  184  (1861).  8  Ibid.,  203  (1861). 


228  THE  NATURE  OF  SOLUTION 

Relative  rates 

of  diffusion 

Sodium  chloride  1.0 

Alcohol  0.476 

Glycerol  0.440 

Mannite  0.349 

Milk-sugar  0.185 

Cane-sugar  0.214 

Gum  arabic  0.004 

These  data  show  the  almost  non-diffusability  of  col- 
loids as  compared  with  crystalloids. 

The  separation  of  crystalloids  from  colloids  by  dif- 
fusion Graham  termed  dialysis,  and  the  apparatus  for 
effecting  such  separations  a  dialyzer.  This  is  one  of  the 
most  efficient  methods  of  freeing  colloids  from  crystalloids. 
The  mixture  is  placed  hi  a  dialyzer,  when  the  crystalloid 
passes  through  and  the  colloid  remains  behind.  We 
shall  see,  however,  that  this  method  of  separating  crys- 
talloids from  colloids  is  often  far  from  quantitative. 
When  colloids  are  precipitated  by  the  addition  of  crystal- 
loids, as  is  often  done,  the  colloid  frequently  carries  down 
some  of  the  crystalloid  with  it,  and  holds  it  so  firmly  that 
it  will  not  diffuse  out  and  leave  the  pure  colloid  behind. 
Frequently,  in  such  cases,  the  colloid  cannot  be  washed 
free  from  the  crystalloid. 

This  method  based  upon  dialysis  is,  however,  a  fairly 
general  method  for  obtaining  colloids  reasonably  free 
from  crystalloid  impurities. 

Graham  gives  a  number  of  examples  of  colloids  pre- 
pared by  dialysis.  Thus,  neutral  aluminium  chloride 
dialyzes  without  undergoing  any  decomposition.  When 
hydrated  alumina  is  dissolved  hi  aluminium  chloride  and 
the  mixture  dialyzed,  there  remains  behind  soluble 
hydrated  alumina.  After  standing  in  the  dialyzer  for 
nearly  a  month  this  was  found  to  contain  only  a  trace  of 
hydrochloric  acid.  Graham  points  out  that  soluble  alu- 
mina is  very  unstable,  being  precipitated  by  a  mere  trace 
of  almost  any  salt  or  acid.  He  obtained  similar  results  with 
chromium  hydroxide. 


COLLOIDAL  SOLUTIONS  229 

Graham  prepared  organic  as  well  as  inorganic  colloids 
by  dialysis.  Thus,  caramel  and  albumen  were  obtained 
in  the  form  of  colloidal  solutions. 

Graham's  Views  on  Colloids.  —  Near  the  close  of  his 
great  paper,  Graham1  summarizes  his  views  on  colloids. 
What  these  were  can  be  seen  best  by  quoting  his  own 
words,  at  least  in  part. 

"I  may  be  allowed  to  advert  again  to  the  radical  dis- 
tinction assumed  in  this  paper  to  exist  between  colloids 
and  crystalloids  in  their  intimate  molecular  constitution. 
Every  physical  and  chemical  property  is  characteristically 
modified  in  each  class.  They  appear  like  different  worlds 
of  matter,  and  give  occasion  to  a  corresponding  division 
of  chemical  science.  The  distinction  between  these  kinds  of 
matter  is  that  subsisting  between  the  material  of  a  mineral 
and  the  material  of  an  organized  mass." 

Again,2  "The  phenomena  of  the  solution  of  a  salt  or 
crystalloid  probably  all  appear  hi  the  solution  of  a  colloid, 
but  greatly  reduced  hi  degree.  The  process  becomes  slow; 
time,  indeed,  appearing  essential  to  all  colloidal  changes. 
The  change  of  temperature,  usually  occurring  hi  the  act  of 
solution,  becomes  barely  perceptible."  .  .  .  "The  col- 
loid, although  often  dissolved  in  a  large  proportion  by 
its  solvent,  is  held  in '"solution  by  a  singularly  feeble  force. 
Hence  colloids  are  generally  displaced  and  precipitated 
by  the  addition  to  their  solution  of  any  substance  from  the 
other  class.  Of  all  the  properties  of  liquid  colloids,  their 
slow  diffusion  hi  water,  and  their  arrest  by  colloidal  septa, 
are  the  most  serviceable  hi  distinguishing  them  from 
crystalloids.  Colloids  have  feeble  chemical  reactions." 
"It3  is  difficult  to  avoid  associating  the  inertness  of  col- 
loids with  their  high  equivalents,  particularly  where  the 
high  number  appears  to  be  attained  by  the  repetition  of  a 
smaller  number.  The  inquiry  suggests  itself  whether  the 
colloid  molecule  may  not  be  constituted  by  the  grouping 

1  Phil.  Trans.,  161,  220  (1861). 

•  Ibid.,  220  (1861).  «  Ibid.,  221  (1861). 


230  THE  NATURE  OF  SOLUTION 

together  of  a  number  of  smaller  crystalloid  molecules,  and 
whether  the  basis  of  colloidality  may  not  really  be  this 
composite  character  of  the  molecule." 

This  remarkable  paper  concludes  with  a  discussion  of 
" Osmose,"1  which  contains  essentially  the  views  that  we 
hold  today.  "It  now  appears  to  me  that  the  water 
movement  in  osmose  is  an  affair  of  hydration  and  of 
dehydration  in  the  substance  of  the  membrane  or  other 
colloid  septum.  The  outer  surface  of  the  membrane  being 
in  contact  with  pure  water  tends  to  hydrate  itself  to  a 
higher  degree  than  the  inner  surface  does,  the  latter  sur- 
face being  supposed  to  be  in  contact  with  a  saline  solu- 
tion." Where  the  membrane  comes  hi  contact  with  the 
solution,  "the  degree  of  hydration  is  lowered,  and  the 
water  must  be  given  up  by  the  inner  layer  of  the  membrane 
and  it  forms  the  osmose." 

This  is  essentially  the  view  we  hold  today  in  connection 
with  the  passage  of  water  through  semipermeable  colloidal 
membranes  in  the  measurement  of  osmotic  pressure. 

A  brief  reference  has,  then,  been  made  to  some  of  the  ex- 
perimental work  of  Graham  on  colloids;  and  to  some  of  his 
more  important  conclusions  in  regard  to  this  state  of  matter. 
All  things  considered,  we  must  regard  this  distinguished 
English  chemist  as  the  father  of  colloidal  chemistry. 

Work  of  M.  Carey  Lea.  —  The  American  chemist  M. 
Carey  Lea2  in  1889  made  a  discovery  which  attracted 
considerable  attention.  He  reduced  a  ten  percent  solu- 
tion of  silver  nitrate  with  a  solution  of  ferrous  sulphate, 
to  which  a  solution  of  sodium  citrate  and  sodium  carbonate 
was  added.  The  precipitated  silver  was  soluble  and  he 
thus  obtained  a  solution  which  contained  more  than  95 
percent  of  silver,  and  such  systems  were  called  by  him 
solutions  of  metallic  silver. 

Although  Loew3  in   1883,  and  Muthmann4   in    1887, 

1  Phil.  Trans.,  151,  223  (1861). 

2  Amer.  Journ.  Sci.,  37,  476  (1889);  38,  47,  237,  241  (1889). 

3  Ber.  d.  chem.  GeselL,  16,  2707  (1883).  4  Ibid.,  20,  983  (1887). 


COLLOIDAL  SOLUTIONS  231 

had  prepared  colloidal  solutions  of  silver,  their  work  at 
that  time  had  not  aroused  any  very  great  interest.  Lea 
regarded  his  preparation  as  a  true  solution  of  metallic 
silver,  or  an  allotropic  modification  of  silver.  When  it  is 
recalled  that  silver  is  one  of  the  most  insoluble  metals  hi 
water,  and  here  was  apparently  a  solution  containing  more 
than  95  percent  of  silver,  we  can  see  the  reason  for  the 
strong  impression,  not  to  say  sensation,  created  by  the 
discovery  of  M.  Carey  Lea. 

In  the  light  of  what  we  know  today,  as  we  shall  see, 
there  is  nothing  very  surprising  hi  this  discovery.  He 
was  dealing  not  with  a  true  solution  of  silver  hi  water,  but 
with  a  colloidal  solution  of  the  metal,  and  we  today  know 
colloidal  solutions  of  the  metals  hi  great  abundance. 
Bredig  and  others  have  prepared  colloidal  solutions  not 
only  of  metals  as  insoluble  as  silver,  but  of  far  more  in- 
soluble metals  such  as  platinum,  iridium,  gold,  and  the 
like. 

This  brings  us  down  to  the  work  which  will  be  dis- 
cussed hi  more  detail  hi  this  book.  The  development  of 
this  new  branch  of  solution  has  made  necessary  the  intro- 
duction and  use  of  certain  terms,  the  meaning  of  which 
must  be  understood  before  the  subject  can  be  considered. 
As  far  as  possible  these  will  now  be  defined. 

Nomenclature,  Terms  Introduced.  —  Several  of  the  terms 
still  in  use  were  introduced  by  Graham  in  his  classical 
researches  on  colloids.  We  have  seen  that  he  distinguished 
between  crystalloids  and  colloids]  the  former  diffusing 
readily  through  membranes  of  vegetable  parchment,  the 
latter  diffusing  only  slowly  or  not  at  all  through  such 
membranes.  The  diffusion  through  the  membrane  he 
called  osmosis.  The  separation  of  crystalloids  from  col- 
loids by  means  of  diffusion  through  such  membranes, 
Graham  termed  dialysis;  and  the  apparatus  hi  which  such 
separations  were  effected,  a  dialyzer. 

He  found  that  many  substances  which  are  ordinarily 
insoluble  hi  water,  such  as  arsenic  trisulphide,  silicic 


232  THE  NATURE  OF  SOLUTION 

acid,  etc.,  when  prepared  in  certain  ways,  apparently 
dissolve  in  water.  Thus,  if  hydrogen  sulphide  is  added 
to  a  solution  of  arsenic  chloride,  the  arsenic  sulphide  is 
precipitated  as  a  yellow  solid.  If,  on  the  other  hand,  an 
aqueous  solution  of  hydrogen  sulphide  is  added  to  an 
aqueous  solution  of  arsenic  oxide,  no  precipitate  is  formed; 
the  arsenic  sulphide  thus  produced  remaining  in  "solution, 
giving  a  yellowish  color  to  the  solution.  Similarly,  if 
hydrochloric  acid  is  added  to  a  strong  solution  of  a  soluble 
silicate,  silicic  acid  Is  precipitated.  If,  however,  hydro- 
chloric acid  in  excess  is  added  to  a  dilute  solution  of  a 
soluble  silicate,  there  is  no  precipitation  of  silicic  acid, 
but  the  solution  remains  perfectly  clear,  indicating  that 
the  silicic  acid  formed  remains  in  solution.  Graham 
dialyzed  such  apparent  solutions  and  found  that  the  silicic 
acid  was  not  present  as  a  true  solution,  but  was  there  as  a 
colloidal  solution. 

When  both  of  the  constituents,  solvent  and  substance, 
of  a  colloidal  solution  were  liquid,  Graham  termed  the 
system  a  sol.  If  one  of  the  constituents  was  solid,  the 
system  appearing  to  be  like  a  jelly,  he  termed  this  a  gel. 

Graham  found  that  liquids  other  than  water  are  capable 
of  forming  colloidal  solutions.  Thus,  alcohol,  glycerol, 
etc.,  could  form  such  solutions.  Colloidal  solutions  hi 
water  he  termed  hydrosols  and  hydrogels',  in  alcohol,  alcosols 
and  alcogels;  in  glycerol,  glycerosok  and  glycerogek.  So 
much  for  the  nomenclature  of  colloidal  chemistry  as  far 
as  we  owe  it  to  Graham. 

More  Recent  Terms.  —  As  our  knowledge  of  colloids 
has  increased  in  recent  times,  it  has  been  found  necessary 
to  supplement  Graham's  nomenclature  with  certain  other 
terms.  We  shall  see  that  the  differences  between  true 
solutions,  colloidal  solutions,  and  colloidal  suspensions,  are 
vitally  connected  with  the  state  of  division  of  the  dissolved 
substance.  In  a  true  solution  the  dissolved  substance,  if  a 
non-electrolyte,  is,  as  we  have  seen,  usually  in  the  molec- 
ular condition  as  shown  by  the  determination  of  its 


COLLOIDAL  SOLUTIONS  233 

molecular  weight  by  the  freezing-point  or  the  boiling-point 
method.  In  colloidal  solutions  the  particles  are,  as  we 
shall  see,  much  more  coarse-grained.  In  colloidal  suspen- 
sions they  are  still  more  coarse-grained,  and  hi  mechanical 
suspensions  more  coarse-grained  than  hi  colloidal  suspen- 
sions. The  degree  of  fine-grainedness  or  dispersity  of  the 
colloidal  particles  is  then  a  matter  of  great  importance  hi 
the  chemistry  of  colloids. 

If  the  dispersion  is  not  great  we  have  emulsions  and 
suspensions.  If  the  dispersion  is  great  the  dispersoids 
are  divided  into  emulsoids  and  suspensoids,  and  these 
terms  are  frequently  used  in  discussing  colloids. 

Nomenclature  of  Wolfgang  Ostwald.  —  Wolfgang  Ost- 
wald  would  divide  dispersoids  as  follows,  depending  on 
then'  degree  of  dispersity  or  fine-grainedness:  (a)  coarse 
dispersions  (these  include  the  emulsions  and  suspen- 
sions), (6)  colloidal  solutions  proper  (finer-grained),  (c) 
molecular  dispersoids  (still  finer-grained),  (d)  ion  disper- 
soids (finest  of  them  all).  Group  6,  or  colloidal  solutions, 
includes  emulsoids  and  suspensoids.  The  emulsoids,  or 
emulsion  colloids,  hi  which  the  dissolved  substance  is  liquid, 
show  certain  fairly  characteristic  properties.  Among  these 
are  to  be  mentioned  their  high  viscosity;  under  the  ultra- 
microscope  the  individual  particles  can  not,  hi  general, 
be  seen,  there  being  only  a  general  illumination  of  the 
field;  very  concentrated  solutions  of  salts  are  required 
to  coagulate  them,  and  they  are  devoid  of  the  interesting 
and  important  electrical  properties,  including  the  carrying 
of  a  charge,  which  are  characteristic  of  the  suspensoids. 

The  suspensoids  or  suspension  colloids  have  also  more 
or  less  well-defined  and  characteristic  properties;  their  vis- 
cosity is  not  very  different  from  that  of  the  pure  solvent; 
when  examined  under  the  ultramicroscope  the  indi- 
vidual light  discs  can  be  seen;  they  are  readily  coagu- 
lated by  electrolytes  and  by  whirling  in  a  centrifuge;  the 
parts  are  charged  electrically,  which  we  shall  see  is  very 
important  hi  determining  the  nature  of  colloidal  solutions, 


234  THE  NATURE  OF  SOLUTION 

as  well  as  in  the  precipitation  of  the  colloidal  particles 
by  electrolytes,  and  they  show  the  Brownian  movement. 
Good  examples  of  the  suspension  colloids  are  the  metal 
hydrasols,  while  gelatine  solutions  are  examples  of  emulsoids. 

While  these  groups  of  dispersoids  are  fairly  distinct 
from  one  another,  yet  systems  are  known  which  corre- 
spond to  transitions  from  the  one  to  the  other.  A  sol  may 
have  the  properties  of  a  gel,  and  vice  versa.  This  apparent 
abnormality  has  been  explained  as  due  probably  to  attrac- 
tion between  the  two  phases  present.  Where  such  attrac- 
tion exists  the  system  has  been  called  a  lyophile;  where 
it  does  not  exist  a  lyophobe.  We  can  have  lyophiles  and 
lyophobes  in  water,  in  alcohol,  in  glycerol,  etc.  In  these 
cases  they  become  hydrophiles  and  hydrophobes,  alcophiks 
and  alcophobes,  glycerophiks  and  glycerophobes,  etc. 

As  we  have  seen,  we  have  both  sols  and  gels,  and  each 
can  pass  into  the  other.  When  a  sol  passes  into  a  gel,  we 
speak  of  it  as  gelation.  When  a  gel  passes  into  a  sol,  we 
refer  to  it  as  solation. 

With  these  terms  clearly  in  mind  we  can  now  proceed 
to  the  study  of  the  more  recent  work  in  the  field  of  col- 
loid chemistry,  and  shall  take  up  first  the  methods  which 
have  been  discovered  and  are  used  for  preparing  colloidal 
solutions  in  the  broad  sense  of  that  term. 

Methods  of  Preparing  Colloidal  Solutions.  —  A  fairly 
large  number  of  methods  have  been  devised  for  preparing 
colloidal  solutions.  These  admit  of  easy  classification,  in 
that  they  fall  into  two  general  classes — Chemical  Methods 
and  Electrical  Methods. 

The  chemical  methods  include  double  decompositions  — 
hydrolysis  being  a  double  decomposition  in  which  water 
is  one  of  the  substances  involved.  Reductions  are  ob- 
viously chemical  methods. 

The  electrical  method,  as  we  shall  see,  has  come  prom- 
inently into  play  in  recent  times. 

Double  Decompositions.  —  We  have  seen  how  Graham 
prepared  colloidal  silicic  acid  and  colloidal  arsenic  sulphide. 


COLLOIDAL  SOLUTIONS  235 

These  are  simply  examples  of  double  decompositions 
which  are  so  well  known  in  chemistry.  Other  examples  of 
methods  involving  simple  double  decomposition  in  which 
there  is  neither  oxidation  nor  reduction,  are  the  following. 
It  is  well  known  that  whenever  the  salt  of  a  weak 
acid  or  a  weak  base,  or  still  better  if  the  acid  and  base 
are  both  weak,  is  brought  into  the  presence  of  water, 
the  salt  is  in  part  hydrolyzed  or  broken  down  into  the  free 
acid  and  the  free  base.  A  fairly  large  number  of  bases 
have,  in  this  way,  been  prepared  in  the  colloidal  condition. 

Although  hydrochloric  acid  is  a  strong  acid,  its  salts 
with  weak  bases  such  as  aluminium  chloride  and  ferric 
chloride  are  more  or  less  hydrolyzed  by  water,  espe- 
cially at  more  elevated  temperatures,  and  the  resulting 
hydroxides  are  in  the  colloidal  condition. 

Some  nitrates  are  also  hydrolyzed,  and  the  resulting 
hydroxides  are  in  the  colloidal  condition.  If  the  acid  is 
weak  as  well  as  the  base  the  hydrolysis  is  greatly  increased. 
Acetates  are  much  more  hydrolyzed  than  chlorides  and 
nitrates,  and  the  hydrolytic  action  of  water  on  acetates 
has  been  used  fairly  extensively  in  the  preparation  of 
colloidal  solutions  of  hydroxides.  In  this  way  colloidal 
solutions  of  iron  and  aluminium  and  other  hydroxides  have 
been  prepared. 

Reductions.  —  So  much  for  the  methods  involving 
simple  chemical  double  decompositions.  Other  chemical 
reactions  have  been  used  to  prepare  colloidal  solutions. 
These  involve  reduction  processes,  and  will  be  briefly 
considered.  The  reactions  have  to  do  primarily  with  the 
reduction  of  salts  of  the  metals  by  certain  reducing  agents. 

These  reduction  methods  were  the  first  to  be  used  hi 
preparing  colloidal  solutions.  Thus,  at  the  very  beginning 
of  the  nineteenth  century  gold  sols  were  prepared  by  re- 
ducing salts  of  gold  with  mild  reducing  agents. 

Faraday,1  hi  1857,  prepared  gold  sols  by  reducing  auric 
chloride  with  yellow  phosphorus  dissolved  in  ether,  and 

1  Phil.  Trans.,  147,  145  (1857). 


236  THE  NATURE  OF  SOLUTION 

the  soluble  silver  of  M.  Carey.  Lea,1  already  referred  to,  was 
prepared  by  reducing  a  silver  salt  with  a  concentrated 
solution  of  ferrous  citrate. 

The  history  of  the  preparation  of  purple  of  Cassius 
is  interesting.  As  is  well  known,  it  is  formed  when  a 
dilute  solution  of  a  salt  of  gold  is  treated  with  stannous 
chloride. 

The  nature  of  this  purple  of  Cassius  has  attracted  the 
attention  of  chemists  since  the  time  of  Berzelius,  but  only 
since  colloidal  chemistry  came  to  the  front  has  its  real 
nature  been  understood.  Gold  sols  have  recently  been 
prepared  by  reducing  salts  of  gold  with  carbon  monoxide. 
Zsigmondy  has  found  formaldehyde  a  satisfactory  reducing 
agent  for  preparing  sols  of  gold  from  gold  salts,  while 
Gutbier  has  made  use  of  the  mild  reducing  agents  hydroxyl- 
amine,  phenylhydrazine,  and  hydrazine.  He  has  also 
prepared  by  this  method  sols  of  the  very  resistant  metals, 
platinum,  iridium,  etc.  The  Gutbier  method  gives  col- 
loidal solutions  of  the  metals  which  are  relatively  stable, 
and  when  properly  protected,  persist  for  considerable 
periods  of  tune. 

Stable  sols  of  platinum  have  also  been  prepared  by 
reducing  its  salts  with  acrolein. 

A  few  colloidal  solutions  have  been  made  by  the  oppo- 
site of  reduction  methods,  viz.,  oxidation  methods;  but 
these  are  not  sufficiently  important  to  call  for  any  detailed 
discussion  here. 

Electrical  Methods.  —  It  was  known  to  Humphrey 
Davy  that  when  a  heavy  electrical  discharge  is  passed 
between  metal  poles,  the  metal  particles  are  torn  off  in 
a  fine  state  of  division. 

Utilizing  this  fact,  Bredig2  has  devised  a  method  for 
preparing  colloidal  solutions  of  many  of  the  most  resistant 
metals.  If  it  is  desired  to  prepare  a  platinum  sol^two 
stout  wires  of  platinum  are  dipped  into  pure  water, 

1  Amer.  Journ.  Science,  37,  476;  38,  47  (1889). 

2  Zeit.  phys.  Chem.,  31,  258  (1899). 


COLLOIDAL  SOLUTIONS  237 

and  the  terminals  brought  sufficiently  close  together  so 
that  an  electric  arc  is  set  up  between  them  when  the 
proper  electromotive  force  is  impressed  upon  the  metal 
wires.  Under  these  conditions  the  metal  in  question  is 
torn  off  in  a  very  fine  state  of  division,  and  remains  sus- 
pended hi  the  water  as  a  sol  of  the  metal  in  question.  In 
this  way  Bredig,  using  a  current  from  30  to  110  volts,  and 
from  5  to  10  amperes,  was  able  to  prepare  sols  of  plati- 
num, iridium,  silver,  and  gold. 

The  method  of  Bredig  worked  very  well  hi  water,  but 
when  other  liquids,  especially  the  organic,  were  employed, 
it  did  not  give  such  satisfactory  results.  The  liquids 
themselves  often  underwent  decomposition,  liberating  car- 
bon which  was  admixed  with  the  sol  in  question. 

The  Bredig  method,  as  modified  by  Svedberg,1  over- 
came some  of  these  difficulties,  and  enabled  sols  to  be 
prepared  hi  solvents  other  than  water. 

The  surface  of  the  metal  of  which  it  was  desired  to 
prepare  a  sol  was  increased.  A  condenser  was  introduced 
into  the  circuit  and  a  very  small  current  of  50  or  less 
milliamperes  was  employed.  Later  Svedberg  used  an 
inductorium  which  gave  an  alternating  current.  With 
this  improved  method  he  prepared  sols  of  silver,  tin,  gold, 
etc.,  in  several  solvents. 

He  prepared  sols  of  all  of  the  alkali  metals  hi  ethyl 
ether  as  well  as  a  large  number  of  metal  sols  in  isobutyl 
alcohol.  He  obtained  sols  not  only  of  most  of  the  metals, 
but  of  selenium,  carbon,  silicon,  etc.,  and  of  a  number  of 
minerals. 

PROPERTIES  OF  COLLOIDAL  SOLUTIONS 

The  three  most  characteristic  properties  of  true  solutions 
are:  osmotic  pressure,  lowering  of  freezing-point,  and  rise 
in  boiling-point.  These  properties  are  not  only  possessed 
by  all  true  solutions,  but  they  obey  certain  well-defined 

1  Ber.  d.  chem.  Gesett.,  38,  3616  (1905);  39,  1705  (1906).  Kott.  Zeit., 
1,229,257(1907). 


238  THE  NATURE  OF  SOLUTION 

laws.  Thus,  the  osmotic  pressures  of  true  solutions  obey 
the  laws  of  gases,  as  we  have  already  seen.  The  lowering 
of  the  freezing-point  of  the  solvent,  and  the  rise  hi  its 
boiling-point  produced  by  dissolved  substances,  are  in 
accord  with  the  well-known  laws  of  Raoult  dealing  with 
these  phenomena.  Given  a  system  which  appears  to  be  a 
true  solution,  the  question  is,  how  can  we  determine 
whether  or  not  it  is  a  true  solution?  Knowing  its  concen- 
tration, we  measure  its  osmotic  pressure,  its  freezing  and 
boiling-point,  and  see  whether  all  of  these  conform  to  the 
well-known  laws  of  true  solutions. 

uOne  of  the  first  questions  that  would  naturally  arise  in 
connection  with  colloidal  solutions  is,  do  they  have  these 
three  properties  at  all,  and  if  so,  to  what  extent? 

Osmotic  Pressures  of  Emulsions  and  of  Suspensions. — 
The  above  question  resolves  itself  into  two.  Do  emulsoids 
or  emulsion  colloids  have  osmotic  pressure,  and  if  so  to 
what  extent?  Do  suspensoids  show  osmotic  pressure,  and  if 
so  what  is  its  order  of  magnitude?  Take  first  the  emul- 
sions. Pfeffer  measured  the  osmotic  pressures  of  a  few  col- 
loids and  found  values  which  are  very  small  as  compared 
with  crystalloids. 

Rodewald  and  Kattein1  measured  the  osmotic  pressures 
of  the  starch  iodine  emulsoid.  A  sol  containing  about  30 
grams  to  the  liter  gave  an  osmotic  pressure  of  approxi- 
mately 20  cm.  of  water. 

The  most  elaborate  work  in  this  field  is  perhaps  that 
of  Lillie,2  who  studied  the  osmotic  pressures  of  a  number 
of  emulsoids.  He  found  that  a  solution  of  egg  albumin 
containing  12.5  grams  to  the  liter  showed,  at  ordinary 
temperatures,  an  osmotic  pressure  of  about  20  mm.  of 
mercury.  When  the  osmometer  reached  this  height  it 
stood  constant  for  a  considerable  period  of  time.  This 
fact  is  of  importance  as  bearing  on  the  question  as  to 
whether  the  osmotic  pressure  actually  observed  was  due 
to  the  emulsoid  itself,  or  to  some  crystalloid  impurity 

1  Zeit.  phys.  Chem.,  33,588  (1900).    2  Amer.  Journ.  PhysioL,  20, 127  (1907). 


COLLOIDAL  SOLUTIONS  239 

mixed  with  it.  If  it  were  due  to  some  crystalloid  mixed 
with  the  colloid,  then  on  standing  in  contact  with  the 
membrane  this  would  have  diffused  out  of  the  colloid 
through  the  membrane,  and  the  pressure  registered  on  the 
osmometer  would  not  have  remained  constant  over  any 
considerable  period  of  time.  The  fact  that  this  pressure 
did  remain  constant  indicates  that  it  was  due  to  the  emul- 
soid  itself.  The  only  alternative  seems  to  be  to  assume 
that  it  is  due  to  some  admixed  crystalloid  which  is  so 
held  by  the  colloid  that  it  cannot  diffuse  out  of  it  through 
the  membrane.  This  alternative  seems  to  be  scarcely 
acceptable.  When  we  take  into  account  the  further  fact 
that  an  emulsoid  prepared  in  a  number  of  different  ways, 
and  therefore,  which  could  not  contain  even  the  same 
crystalloid,  much  less  the  same  or  comparable  amounts  of 
crystalloids,  shows  the  same  osmotic  pressure,  we  are  almost 
forced  to  the  conclusion  that  the  osmotic  pressure  hi  question 
is  due  to  the  emulsoid,  and  not  to  some  admixed  impurity. 

When  we  turn  to  suspensoids  the  results  of  the  direct 
measurements  of  osmotic  pressure  are  less  satisfactory. 
Linder  and  Picton1  worked  with  sols  of  ferric  hydroxide 
and  arsenic  trisulphide,  attempting  to  measure  directly 
their  osmotic  pressures.  They  were  not  able  to  obtain 
concordant  results.  If  these  systems  have  osmotic  pres- 
sures it  can  be  said  that  these  pressures  are  very  small. 
This  may  be  due  in  part  to  the  relatively  large  sizes  of 
the  particles  in  a  colloidal  suspension. 

Lowering  of  Freezing-Point  and  of  Vapor  Tension.  —  The 
freezing-point  lowerings  and  lowerings  of  vapor-tension  or 
rise  in  boiling-point,  produced  even  by  emulsoids,  must  of 
necessity  be  small,  since  the  osmotic  pressures  are  so  small. 
This  will  be  obvious  when  we  consider  that  an  osmotic 
pressure  of  1  mm.  of  mercury  corresponds  to  only  about 
one  ten-thousandth  of  a  degree  lowering  of  the  freezing- 
point.  Bruni  and  Pappada2  prepared  some  very  pure  sols 

1  Journ.  Chem.  Soc.,  87,  1906  (1905). 

*  Rend.  R.  Ace.  dei  Lined  [5],  9,  354  (1900). 


240  THE  NATURE  OF  SOLUTION 

of  albumin,  gelatine,  etc.,  and  were  unable  to  detect 
any  difference  between  their  freezing-points  and  boiling- 
points,  and  those  of  pure  water.  This  shows  how  mean- 
ingless are  a  number  of  the  attempts  that  were  made 
to  determine  the  molecular  weights  of  colloids  by  the 
freezing-point  and  the  boiling-point  methods.  Enormous 
molecular  weights  were  found  for  many  colloids,  and  then* 
molecular  weights  may  be  relatively  enormous,  but 
there  is  no  necessary  connection  between  the  molecular 
weights  found  by  these  methods  and  the  true  molecular 
weights  of  these  substances.  It  is  highly  probable  that 
in  most  cases  the  lowering  of  the  freezing-point  and  the 
rise  in  the  boiling-point  observed,  were  due  to  crystalloid 
impurities  hi  the  colloids.  Further,  if  the  colloids  them- 
selves actually  produced  the  phenomena  observed,  it  is 
not  at  all  certain  that  these  properties  could  be  used 
to  determine  the  molecular  weights  of  these  colloids. 
Before  this  can  be  done  it  must  be  shown  that  Raoult's 
laws  of  freezing-point  lowering  and  depression  of  vapor- 
tension,  or  rise  in  boiling-point,  hold  for  colloids  as  well  as 
for  crystalloids. 

If  these  methods  cannot  be  used  to  determine  the 
molecular  weights  of  emulsoids,  still  less  can  they  be 
employed  with  suspensoids,  which  show  still  smaller  depres- 
sion of  the  freezing-point  and  rise  in  the  boiling-point,  if 
they  have  these  properties  at  all. 

Diffusion  of  Colloids.  —  If  we  recall  the  relations  al- 
ready pointed  out  when  we  were  considering  true  solu- 
tions, between  diffusion  and  osmotic  pressure,  viz.,  that 
all  diffusion  is  caused  by  osmotic  pressure,  we  should  ex- 
pect to  find  all  colloids  diffusing  very  slowly,  if  diffusing 
at  all.  Such  was  shown  to  be  the  case  by  Graham  in 
his  classical  researches  on  colloids.  Indeed,  his  funda- 
mental method  of  distinguishing  between  crystalloids  and 
colloids  was  based  upon  the  powers  of  these  two  classes 
of  substances  to  diffuse,  or  not  to  diffuse,  through  vege- 
table parchment. 


COLLOIDAL  SOLUTIONS  241 

We  have  seen  that  the  emulsoids  show  greater  osmotic 
pressures  than  the  suspensoids  and  we  should  expect  them 
to  diffuse  more  rapidly.  Such  is  the  fact.  Herzog1  has 
determined  the  relative  rates  of  diffusion  of  a  number  of 
colloids  as  compared  with  certain  organic  crystalloids  and 
a  few  of  his  results  are  given  below. 

Diffusion 
constant 

Urea  1.01 

Glucose  0.57 

Pepsin  0.063 

Albumin  0.054 

Emulsin  0.036 

These  suffice  to  show  the  relative  slowness  with  which 
the  colloids  diffuse,  and  these  colloids  are  emulsoids. 
Since  suspensoids  have  smaller  osmotic  pressures  than 
emulsoids  we  should  expect  them  to  diffuse  even  more 
slowly,  and  here  again  the  facts  are  hi  accord  with  pre- 
diction. Indeed,  some  experiments  have  been  carried  out 
which  would  seem  to  make  it  doubtful  whether  certain 
suspensoids,  at  least  under  certain  conditions,  diffuse  at  all. 

Brownian  Movement. —  The  English  botanist  R.  Brown2 
as  early  as  1827  observed  that  the  pollen  grains  of  plants, 
when  suspended  in  water,  were  continually  hi  a  state  of 
oscillatory  motion.  He  observed  that  this  was  a  general 
phenomenon.  Very  finely  divided  solid  particles  suspended 
hi  any  liquid  which  was  not  too  viscous,  showed  these 
movements.  From  the  name  of  the  discoverer  this  has 
come  to  be  known  as  the  Brownian  movement. 

It  was  at  first  thought  that  these  movements  might  be 
due  to  some  cause  external  to  the  liquid,  such  as  jarring, 
changes  hi  temperature  producing  currents  in  the  liquid, 
etc.  This  was  tested  by  working  hi  places  free  from 
mechanical  disturbance  and  at  a  constant  temperature. 
The  movements  of  the  particles  under  these  conditions 
were  essentially  the  same  as  when  less  precautions  were 

1  Zeit.  Elektrochem.,  13,  533  (1907). 
a  Pogg.  Ann.  14,  294  (1828). 


242  THE  NATURE  OF  SOLUTION 

taken,  as  was  shown  by  Wiener.1  The  persistence  of  the 
Brownian  movement  unchanged  for  apparently  unlimited 
periods  of  time,  shows  that  it  could  not  be  due  to  external 
conditions  which  are  constantly  changing. 

Distance  Traveled  and  Velocities  of  the  Particles.  — 
The  distance  through  which  the  particles  move  depends 
largely  upon  their  size.  Particles  having  a  diameter  of 
about  1  //,  oscillate  along  a  path  which  has  a  length  of 
about  2  /z;  smaller  particles,  as  we  shall  see,  move  through 
much  longer  paths.  This  will  be  discussed  under  the  re- 
sults of  work  with  the  ultramicroscope. 

The  velocities  with  which  the  particles  move  are  also  a 
function  of  their  size.  Exner2  showed  that  particles  rang- 
ing hi  diameter  from  0.4  to  1.3  jit  move  with  a  velocity 
of  from  3.8  to  2.7 /z  per  second.  Very  much  smaller 
particles  as  seen  under  the  ultramicroscope  move  with 
velocities  as  much  as  thirty  times  greater.  This  also  will 
be  discussed  when  the  results  of  ultramicroscopic  investiga- 
tions are  considered. 

Not  only  the  size  of  the  suspended  parts  determines 
the  lengths  of  the  paths  they  travel  and  the  velocities  with 
which  they  move,  but,  as  we  would  expect,  the  nature 
of  the  medium  hi  which  the  parts  are  suspended  exerts 
a  pronounced  influence. 

Svedberg3  has  established  two  generalizations  connect- 
ing the  nature  of  the  medium  with  the  motions  of  the 
particles  suspended  in  it.  The  first  connects  the  ampli- 
tude of  the  vibrations  with  the  viscosities  of  the  media. 
The  generalization  hi  question  is4  for  particles  of  a  given 
size,  "if  the  amplitudes  (of  vibrations  of  the  particles)  are 
plotted  as  ordinates,  and  the  viscosities  of  the  media 
as  abscissae,  the  curve  takes  the  form  of  a  hyperbola." 
This  means  that  the  amplitude  is  inversely  proportional 
to  the  viscosity. 

The  other  relation  has  to  do  with  the  time  required  for  a 

1  Pogg.  Ann.,  118,  79  (1863).          8  Zett.  Elektrochem.,  12, 853, 909  (1)906. 
8  Ann.  d.  Phys.  2,  843  (1900).         4  Ibid.,  854  (1906). 


COLLOIDAL  SOLUTIONS  243 

particle  to  move  along  its  path  and  the  velocity  with  which 
it  travels.  "The  velocity1  hi  solvents  of  very  different  na- 
ture is  nearly  constant  and  has  the  value  2  to  4xlO~2 
centimeters  per  second." 

The  tune  of  oscillation  is  approximately  proportional 
to  the  amplitude.  Temperature  has  a  marked  effect. 
Exner2  has  shown  that  between  the  temperatures  20°  and 
70°,  the  square  of  the  velocities  is  proportional  to  the  tem- 
perature. 

Cause  of  the  Brownian  Movement.  —  So  much  for 
some  of  the  facts  pertaining  to  the  Brownian  movement. 
The  important  question  still  remains,  what  causes  these 
movements?  We  have  seen  that  they  cannot  be  due  to 
such  external  causes  as  jarring,  changes  hi  temperature, 
etc.  Since  these  movements  are  not  due  to  external 
causes,  their  cause  must  be  sought  for  within  the  liquid 
itself.  In  the  opinion  of  Ramsay,3  Einstein4  and  others, 
the  Brownian  movement  is  due  to  the  impacts  or  blows 
of  the  molecules  of  the  liquid  against  the  solid  suspended 
particles.  From  the  mathematics  of  chance,  we  might 
naturally  think  that  as  many  liquid  molecules  would 
strike  the  particle  hi  any  given  unit  of  tune  on  the  one 
side  as  on  the  other,  and  the  result  would  be  that  the  par- 
ticle would  remain  at  rest.  Such  would  probably  be  the 
case  if  we  think  of  a  long  period;  but  for  a  relatively  short 
period  a  particle  would  be  hit  on  one  side  by  a  larger 
number  of  molecules  than  on  the  other,  and,  consequently, 
motion  would  result. 

Work  of  Perrin.  —  The  theory  of  the  Brownian  move- 
ment has  been  worked  out  by  Einstein  and  Smoluchowski. 
It  has  been  studied  experimentally  by  Ehrenhoft,  Sved- 
berg,  and  by  Perrin.  The  work  of  the  last  named  has 
attracted  so  much  attention  that  one  phase  of  it  will  be 
considered  in  some  detail.  Perrin  undertook  his  work  to 
determine  the  cause  of  the  Brownian  movement,  and  to 

1  Zett.  Elektrochem.,  12,859.  3  Chem.  News,  66,  90  (1892). 

2  Ann.  d.  Phys.,  2,  843  (1900).  *  Ann.  d.  Phys.,  19,  371  (1906). 


244  THE  NATURE  OF  SOLUTION 

find  out  whether  an  emulsion  obeyed  the  gas  laws  as  true 
solutions  do.  He  reasoned  that  if  large  molecules  such  as 
cane  sugar,  quinine  sulphate  and  the  like,  in  true  solutions, 
obeyed  the  gas  laws  as  well  as  molecules  of  smaller  molec- 
ular weights,  then  there  seemed  to  be  no  a  priori  reason 
why  the  grains  of  an  emulsion  should  not  obey  these  same 
laws. 

If  an  emulsion  obeys  the  gas  laws,  and  we  allow  a 
uniform  emulsion  to  come  to  equilibrium  at  constant  tem- 
perature, the  grains  ought  to  distribute  themselves  in  the 
emulsion  according  to  the  height,  just  as  in  the  case  of  a 
column  of  gas  or  in  the  atmosphere. 

On  the  assumption  that  the  gas  laws  hold  for  emulsions, 
Perrin  deduced  the  following  equations,1 

2.30  W  log  ^  =  27ra3(A  -  d)  gh 
n 

in  which  W  is  the  mean  granular  energy,  n  and  n0  the  con- 
centrations at  the  two  different  heights,  a  the  radius  of 
the  grains,  A  the  density  of  the  grains,  d  the  density  of 
the  liquid,  g  the  pull  of  gravity  and  h  the  difference 
between  the  two  heights. 

o  r>/77 

From  the  gas  laws  W  =  —r.     Substituting,2  we  have 


2.30        log      =  1  Tra3  (A  -  d)  hg 

which  enables  us  to  calculate  the  Avogadro  constant  N. 
Perrin  prepared  emulsions  hi  water  of  gamboge  and 
of  mastic,  and  by  fractional  centrifuging  he  was  able  to 
obtain  them  practically  uniform.  He  determined  the 
density  of  the  grains  by  two  methods.  First,  he  dried 
the  grains  to  constant  weight  at  about  110°,  and  then 
heated  them  higher,  when  a  viscous  liquid  resulted,  which, 
on  cooling,  formed  a  transparent  glass.  The  density  of 
this  was  obtained  by  suspending  in  a  solution  of  potassium 

1  Ann.  Chim.  Phys.  [8}  18,  32-36  (1909).  2  Ibid.,  p.  62. 


COLLOIDAL  SOLUTIONS  245 

bromide  of  known  concentration.  Secondly,  at  a  given 
temperature  he  obtained  the  mass  m  of  water,  and  ra' 
of  emulsion  filling  an  equal  volume.  By  drying,  he  found 
the  mass  n  of  the  resin  contained  hi  m'  of  emulsion;  the 
mass  of  the  intergranular  water  being  m'  —  n.  If  d  is  the 
absolute  density  of  water,  the  volume  of  water  having  a 

mass   m=  —  ;    and  that  of  intergranular  water  =  —  ;  —  . 
d  d 

f*y\  tyY}  *      _         M 

The  difference,  —  --  :  —  ,  is  the  volume  of  the  grains,  and 
d          d 

n  divided  by  this  volume,  their  density.  The  values 
obtained  by  these  two  methods  are  concordant. 

Concentrations  of  the  Colloids  at  Different  Levels.  — 
To  obtain  n  and  n0,  Perrin  used  a  small  cell  jV  nnn. 
deep,  into  which  was  introduced  a  drop  of  the  emulsion. 
This  was  covered  with  glass  to  prevent  evaporation.  It 
was  placed  under  a  microscope;  the  objective,  while 
having  great  magnifying  power,  had  only  a  small  depth 
of  field,  so  that  only  the  granules  hi  a  layer  yfo  mm. 
in  thickness  could  be  clearly  seen.  By  raising  and  lower- 
ing the  microscope,  grains  at  different  levels  could  be 
observed  and  the  height  read  on  a  micrometer  screw.  Of 
the  larger  grains  Perrin  was  able  to  obtain  photographs 
at  the  different  levels,  but  with  the  smaller  he  was  com- 
pelled so  to  reduce  the  field  that  only  a  few  grains  could 
be  seen  at  any  one  tune.  The  ratio  between  the  mean 
numbers  of  grams  seen  at  the  two  heights  would,  of  course, 

give  the  ratio  of  the  two  concentrations,  —  . 

n 

Radii  of  the  Particles.  —  Perrin  obtained  the  radii  of 
the  particles,  a,  by  three  different  methods. 

First,  he  applied  the  law  of  Stokes  for  bodies  falling  in 
a  liquid. 

—d)g 


in  which  T\  is  the  viscosity  of  the  medium,  v,  the  velocity 
of  fall,  and  A  and  d  the  densities  of  grains  and  medium 


246  THE  NATURE  OF  SOLUTION 

respectively.  The  rate  at  which  the  particles  fell  or 
settled  he  determined  as  follows.  He  placed  some  of  the 
emulsion  hi  a  capillary  tube,  kept  the  temperature  con- 
stant, and  noted  the  rate  at  which  the  emulsion  cleared. 
The  other  factors  are  known  and  a  could  be  calculated. 

Secondly,  Perrin  counted  the  number  of  grains  hi  a 
known  volume  of  emulsion,  which  gives  the  mass  of  a 
grain,  and,  the  density  being  already  known,  also  the 
radius.  The  counting  was  made  possible  by  the  fact  that 
if  the  emulsion  is  rendered  slightly  acid,  the  grains,  when 
they  come  close  to  the  wall,  adhere  to  it,  and  being  at 
rest,  can  be  counted. 

The  third  method  is  based  upon  the  fact  that  when  a 
drop  of  the  emulsion  is  allowed  to  evaporate  on  a  microm- 
eter objective,  the  grains  arrange  themselves  in  lines  which 
can  be  measured;  and  by  dividing  by  the  number  in  the 
line  we  obtain  the  diameter  and  therefore  a. 

These  three  methods  gave  concordant  results,  showing 
that  the  law  of  Stokes  can  be  applied  to  emulsions. 

Perrin  found  that  it  was  a  comparatively  simple  matter 
to  verify  the  law  that  at  equal  elevations  there  were 
equal  rarefactions.  He  used  grains  whose  radii  varied  as 
much  as  one  to  forty.  He  used  different  resins,  varying 
the  viscosity,  but  always  obtained  concordant  results. 
He  found  for  the  value  of  n,  70.5  X 1022,  and  for  the  weight 
of  the  hydrogen  atom  in  one  case  1.47x10- 24  grams,  which 
agrees  closely  with  the  value  1.6x10  ~24,  which  was  derived 
from  the  equation  of  Clausius  and  Maxwell.  This  is 
good  evidence  that  the  Brownian  movement  is  to  be 
explained  by  the  kinetic  theory. 

Mass  of  the  Atom.  —  Perrin  studied  the  Brownian 
movement  to  determine  the  mass  of  the  atom.  Using  the 
equation  of  Einstein,  which  involves  the  mean  displacement 
of  the  grain  in  a  given  time  and  the  mean  energy  of 
rotation,  he  obtained  for  the  weight  of  the  hydrogen  atom 
the  values  1 .45  X 10  " 24,  and  1 .56  X 10  " 24.  Perrin  summarizes 
his  work  and  results  as  follows: 


COLLOIDAL  SOLUTIONS  247 

(1)  The  preparation  of  uniform  emulsions,  chosen  at 
will  and  measured  exactly. 

(2)  Extension   of   the   law   of   Stokes   to   microscopic 
dimensions. 

(3)  Demonstration  that  the  gas  laws  apply  to  uniform 
emulsions. 

(4)  Determination  of  the  weights  of  atoms  and  mole- 
cules. 

(5)  Experimental    confirmation    of    the    theories    of 
Einstein. 

THE  ULTRAMICROSCOPE 

We  have  learned  that  by  means  of  the  microscope  we 
can  see  the  small  particles  hi  many  emulsions  and  observe 
their  Brownian  movement.  In  many  emulsions  the  par- 
ticles are  too  small  to  be  seen,  even  with  the  most  power- 
ful microscope.  A  very  important  advance  has  been  made 
in  this  field  by  the  discovery  of  the  ultramicroscope.  We 
owe  this  discovery  primarily  to  Zsigmondy  and  Sieden- 
topf.1  Another  form  of  ultramicroscope  was  developed  hi 
the  same  year  by  Cotton  and  Mouton.2 

Principle  of  the  Ultramicroscope.  —  When  a  beam  of 
light  is  passed  through  dust-free  air,  and  an  attempt  made 
to  observe  it  in  a  direction  at  right  angles  to  the  beam, 
nothing  is  seen.  If  the  air  contains  particles  of  dust,  these 
reflect  the  light  and  the  beam  is  readily  seen. 

Similarly,  when  a  beam  of  light  is  passed  through  a  per- 
fectly clear  solution,  and  an  attempt  made  to  observe  the 
beam  in  a  direction  at  right  angles  to  that  of  propagation, 
nothing  is  seen.  If,  on  the  other  hand,  the  liquid  contains 
suspended  particles,  these  disperse  the  light,  and  the  path 
of  the  beam  through  the  liquid  can  readily  be  seen.  This 
is  what  is  known  as  the  Tyndall  effect.  It  would  seem 
only  natural  to  apply  the  principle  illustrated  by  this 
effect  to  the  study  of  colloidal  solutions;  especially  to  those 
in  which  the  particles  are  too  small  to  be  seen  by  the 

1  Ann.  d.  Phys.,  10,  1  (1903).  *  C&mpt.  Rend.,  136,  1657  (1903). 


248  THE  NATURE  OF  SOLUTION 

microscope;  the  limit  of  microscopic  visibility  being  about 
0.15  M. 

It  was  shown  by  Fizeau1  and  by  Ambronn2  that  beams 
of  light  so  narrow  that  they  fell  below  the  limit  of  the 
microscope,  could  be  seen.  This  suggested  to  Zsigmondy 
that  the  particles  hi  a  colloidal  solution  which  were  too 
small  to  be  seen  by  the  microscope,  might  reflect  enough 
light  to  be  visible. 

The  Zsigmondy  Apparatus.3  —  The  apparatus  at  first  de- 
signed by  Zsigmondy  may  be  roughly  described  as  follows. 
A  beam  of  sunlight  is  reflected  by  a  mirror  through  a  lens, 
which  focuses  it  just  under  the  objective  of  a  microscope, 
in  a  colloidal  solution  contained  in  a  cell  with  glass  walls. 

With  this  apparatus  Zsigmondy,  in  1900,  studied  col- 
loidal solutions  of  gelatine,  glue,  etc.,  and  observed  the 
spots  of  light  reflected  from  the  individual  particles  of  the 
colloid.  In  emulsions  in  which  the  particles  were  very 
small,  the  light  reflected  from  the  individual  particles  could 
not  be  seen. 

Having  found  that  particles  which  are  submicroscopic 
can  be  rendered  visible  by  side  illumination  in  an  other- 
wise dark  field,  Siedentopf  and  Zsigmondy  took  up  the 
question  of  perfecting  the  apparatus,  to  render  visible  still 
smaller  colloidal  particles. 

The  improved  form  of  the  ultramicroscope  of  the  above 
named  investigators  is  based  primarily  upon  two  funda- 
mental principles. 

(a)  The  particles  must  be  illuminated  as  intensely  as 
possible,  care  being  taken  that  none  of  the  light  falls  directly 
upon  the  eye.  Only  the  light  reflected  from  the  colloidal 
particles  passes  up  through  the  microscope  into  the  eye. 

(6)  The  field  of  the  microscope  must  be  kept  dark, 
the  only  light  entering  it  being  from  the  side,  as  already 
discussed. 

i  Pogg.  Ann.,  116,  478  (1862).  •  Wied.  Ann.,  48,  717  (1893). 

8  For  details  see  Colloids  and  the  Ultramicroscope,  by  Zsigmondy;  trans- 
lated by  Alexander  (Longmans). 


COLLOIDAL  SOLUTIONS  249 

The  Siedentopf  Ultramicroscope.  —  The  apparatus  de- 
signed by  Siedentopf1  contains  the  following  essential 
parts.  A  beam  of  sunlight  allowed  to  enter  a  darkened 
room  passes  through  a  telescope  objective  of  focal  length 
about  10  mm.,  and  casts  an  image  of  the  sun  having  a 
diameter  of  about  1  mm.  on  a  slit  which  admits  of  careful 
adjustment.  The  width  of  the  slit  can  be  easily  measured. 
In  the  path  of  the  light  there  is  a  second  telescope  objec- 
tive having  a  focal  length  of  about  80  mm.  The  light 
thus  passes  through  a  microscope  objective  which  projects 
it  into  the  colloidal  solution  contained  in  the  glass  cell. 
It  is  reflected  from  the  colloidal  particles  upward  through 
a  microscope  into  the  eye. 

Nomenclature  of  Ultramicroscopy.  —  The  study  by  means 
of  the  ultramicroscope  of  particles  too  small  to  be  seen  by  the 
ordinary  microscope,  has  made  it  desirable,  indeed  necessary, 
to  extend  the  nomenclature  of  microscopy. 

Siedentopf  terms  a  particle  which  is  too  small  for 
ordinary  microscopic  resolvability,  ultramicroscopic.  This 
term  applies  to  particles  whose  diameter  is  less  than  one- 
fourth  JU. 

If  the  ultramicroscopic  particle  can  be  rendered  visible 
it  is  termed  submicroscopic,  and  such  a  particle  is  a  sub- 
micron. 

If  the  ultramicroscopic  particle  cannot  be  rendered 
visible  it  is  termed  amicroscopic,  and  such  a  particle  an 
amicron. 

Results  Obtained  with  the  Ultramicroscope  —  Sizes  of 
the  Colloidal  Particles.  —  One  of  the  first  problems  in 
using  the  ultramicroscope  is  to  obtain  the  liquid  as  nearly 
as  possible  transparent.  Ordinary  distilled  water  does 
not  fulfill  this  condition.  It  contains  colloidal  particles  or 
dust  particles  which  interfere  with  its  use  in  ultramicro- 
scopy.  In  studying  particles  which  are  so  small  that  they 
are  barely  visible,  it  is  very  desirable  to  obtain  water 
which  shows  no  light-cone.  This  is  accomplished  by 
1  Druckshe.  Verz.  opt.  Werk.,  C.  Zeiss,  164  (1904). 


250  THE  NATURE  OF  SOLUTION 

redistilling  the  water,  and  allowing  it  to  stand  for  some 
time  in  a  space  free  from  dust. 

One  of  the  most  interesting  applications  of  the  ultra- 
microscope  is  to  the  problem  of  the  size  of  colloidal  particles. 
Siedentopf  and  Zsigmondy1  have  described  two  methods 
for  determining  their  sizes. 

If  we  represent  the  mass  of  the  colloidal  substance  in 
solution  by  m,  the  number  of  colloidal  particles  by  n,  and 
the  density  of  the  colloidal  substance  by  d,  the  volume 
of  a  single  particle,  v,  is  calculated  from  the  following 
equation  — 

m 
dn 

m,  the  mass  of  the  substance  in  a  given  volume  of  the 
solution  is  obtained  directly,  knowing  the  concentration. 
The  number  of  particles,  n,  in  this  volume  is  counted  by 
means  of  the  ultramicroscope,  the  density,  d,  of  the  sub- 
stance from  which  the  colloid  is  made  being  determined 
in  the  usual  way.  All  of  the  quantities  in  the  above  equa- 
tion are  known  except  v,  which  is,  of  course,  calculated 
at  once.  The  dispersity  of  a  sol  is  usually  expressed  in 
terms  of  the  diameter,  which  is  calculated  from  the  volume 
on  the  assumption  that  the  particle  is  a  cube  or  a  sphere. 

In  counting  the  number  of  particles  in  a  given  volume, 
the  colloidal  solution  must  be  so  dilute  that  the  number 
in  the  illuminated  field  is  very  small.  If  this  is  not  so, 
the  continuous  and  comparatively  rapid  motion  of  the 
particles  obviously  would  make  it  impossible  to  count 
them.  When  favorable  conditions  have  been  secured,  the 
accuracy  of  this  method  is  placed  by  Siedentopf  at  about 
twenty  percent. 

The  second  method  of  determining  the  size  of  the  par- 
ticles is  based  upon  the  expression  — 

mr3 

M:  "'-':  "  =  ^~  ^;,-.fi^ 

1  Colloids  and  the  Ultramicroscope,  p.  117. 


COLLOIDAL  SOLUTIONS  251 

All  of  the  symbols  have  the  same  significance  as  hi  the 
former  expression;  the  new  symbol  r  being  the  average 
distance  between  the  particles.  The  layer  studied  must  be 
accurately  measured  and  the  mean  of  a  number  of  results 
is  usually  taken.  The  depth  must  be  as  great  or  greater 
than  the  distance  between  the  particles.  The  smallness  of 
the  particles  which  can  be  seen  depends  largely  upon  the 
illumination;  and,  as  we  would  expect,  somewhat  upon  the 
nature  of  the  particles.  When  the  source  of  the  illumina- 
tion is  an  electric  light  the  smallest  gold  particles  that  can 
be  seen  have  a  diameter  of  15  /*/*>  which  is  15xlO~6  mm. 
When  very  bright  sunlight  is  used  particles  can  be  seen 
which  are  as  small  as  5  /x/z. 

As  Zsigmondy  points  out,  the  difference  between  the 
index  of  refraction  of  gold  and  that  of  water  is  very  great, 
which,  of  course,  aids  ultramicroscopic  visibility.  Parti- 
cles of  very  few  substances  less  than  15  MM  in  diameter 
can  be  seen  with  the  ultramicroscope  with  an  electric 
light,  and  perhaps  none  smaller  than  5  /z/*  with  the  aid 
of  sunlight.  This  shows  how  entirely  erroneous  is  the 
view  that  with  the  ultramicroscope  molecules  or  ions  of 
any  substance  can  be  seen. 

Motion  of  Emulsion  Particles.  —  We  have  seen  what 
is  meant  by  the  Brownian  movement  of  suspended  par- 
ticles. The  question  arises,  are  the  particles  hi  an  emulsion 
hi  motion?  If  so,  how  do  they  move?  Zsigmondy  has 
described  what  he  observed.  The  particles  move  very 
rapidly  in  very  irregular  paths;  indeed,  so  rapidly  that  the 
Brownian  movement  of  the  larger  suspended  particles  is 
slow  by  comparison.  The  small  particles  hi  a  gold  hydro- 
sol  show  both  translatory  and  oscillatory  motion.  The 
translatory  motion  is  from  one  hundred  to  one  thousand 
times  the  diameter  of  the  particle.  This  takes  place  hi 
about  one-seventh  of  a  second.  The  oscillatory  motion 
has  a  much  shorter  time  period. 

The  velocities  with  which  the  smaller  gold  particles 
move  are  thus  incomparably  greater  than  the  Brownian 


252  THE  NATURE  OF  SOLUTION 

movements  of  larger  suspended  particles.  The  motion 
of  the  smaller  particle  of  gold  seems  to  differ  from  the 
Brownian  movement  of  larger  suspended  particles  in  the 
character  of  the  paths  which  they  describe,  as  well  as  in 
the  velocities  with  which  they  move. 

We  have  seen  that  the  Brownian  movement  becomes 
slower  and  slower  the  larger  the  suspended  particles. 
Just  so  it  is  with  the  motions  of  the  particles  of  gold  in  a 
gold  hydrosol.  The  motion  becomes  slower  and  slower 
the  larger  the  size  of  the  particles.  Zsigmondy  at  one  time 
seemed  to  think  that  the  motion  of  the  small  gold  particles 
in  a  gold  sol  was  something  quite  different  from  the 
ordinary  Brownian  movement.  He  studied  the  cause  of 
these  movements. 

Under  the  ultramicroscope  the  gold  sol  is  illuminated 
from  the  side.  This  might  produce  unequal  heating  in 
the  sol  and  thus  give  rise  to  motion.  Zsigmondy  passed 
the  light  through  water  before  allowing  it  to  enter  the  sol. 
He  thus  cut  out  the  heat  rays,  but  observed  the  same  mo- 
tions of  the  gold  particles. 

He  went  further.  He  illuminated  the  field  for  dif- 
ferent periods  of  time,  but  this  produced  no  effect  on  the 
motions;  showing  that  the  cause  of  the  motion  must  be 
sought  for  within  rather  than  without  the  colloidal  solu- 
tion. 

Zsigmondy  seems  to  think  that  the  fact  that  the  gold 
particles  are  charged  with  one  sign  and  the  water  about 
them  with  the  other,  has  much  to  do  with  the  motions  of 
these  particles.  This  fact  must  certainly  be  taken  into 
account  in  dealing  with  this  phenomenon. 

On  the  other  hand,  it  may  be  shown  that  the  move- 
ments of  the  fine  particles  hi  an  emulsion  are  little  or 
nothing  else  than  very  rapid  Brownian  movements. 

ELECTRICAL  PROPERTIES  OF  COLLOIDS 

When  we  were  studying  true  solutions  we  saw  that  then* 
electrical  properties  were  among  the  most  interesting  and 


COLLOIDAL  SOLUTIONS  253 

important.  The  question  of  carrying  or  not  carrying  an 
electrical  charge  was  the  distinguishing  feature  between  elec- 
trolytes and  non-electrolytes  in  solution.  By  this  property 
primarily,  all  chemical  compounds  were  divided  into  these 
two  great  classes.  The  cause  of  the  marked  differences 
between  the  two  was  found  hi  the  fact  that  the  electrolytes 
in  the  presence  of  water  and  other  dissociating  solvents, 
break  down  into  charged  parts  or  ions;  the  cations  carry- 
ing positive  charges  and  the  anions  negative  charges. 
The  non-electrolytes  are  not  thus  dissociated,  remaining 
in  solution  in  the  molecular  condition. 

We  have  seen  how  fundamentally  different  the  proper- 
ties of  these  two  great  classes  of  substances  are,  and  how 
these  differences  are  to  be  referred  to  the  charged  ions  on 
the  one  hand,  and  to  the  uncharged  molecules  on  the  other. 

The  question  of  the  colloidal  particles  being  charged  or 
uncharged,  is  quite  as  fundamental  for  colloids  as  for  true 
solutions  of  electrolytes  and  non-electrolytes.  The  charges 
on  the  colloidal  particles  not  only  determine  many  of  their 
physical  properties  but,  as  we  shall  see,  the  stability  of  the 
colloid  itself. 

Electrical  Endosmosis.  —  This  term  is  applied  to  the 
passage  of  a  liquid  through  a  diaphragm,  or  through 
capillary  tubes,  when  a  current  is  passed  between  two  elec- 
trodes placed  the  one  at  each  end  of  the  tube.  Take  a  U-- 
shaped glass  tube1  and  place  at  the  bottom  of  the  U  a 
bundle  of  short  capillary  tubes.  Insert  one  electrode  at 
each  end  of  the  tubes  and  pass  a  current  from  one  to  the 
other  through  the  liquid  which  fills  the  U-tube.  The  liquid 
will  rise  in  one  of  the  arms  of  the  U-tube,  which  one 
depending  on  the  direction  in  which  the  current  is  passed, 
until  it  stands  at  a  higher  level  in  one  arm  than  in  the 
other.  For  a  definite  difference  in  potential  between  the 
electrodes  the  liquid  will  rise  to  a  definite  height  in  the  one 
arm,  establishing  a  definite  difference  in  level  in  the  two 
arms,  when  equilibrium  will  be  reached.  This  phenome- 

1  Freundlich:  Kapillarchemie,  p.  224  (1909). 


254  THE  NATURE  OF  SOLUTION 

non  seems  to  have  been  observed  for  the  first  time  by 
Reuss.1  It  was  subsequently  studied  by  G.  Wiedemann2 
and  by  Quincke,3  and  later  investigated  quantitatively  and 
mathematically  by  Helmholtz.4  Perrin6  has  quite  recently 
worked  out  the  theory  of  the  process. 

The  height  to  which  the  liquid  will  rise,  or  the  difference 
hi  level  of  the  liquid  hi  the  two  arms  of  the  U-tube  already 
referred  to,  is,  as  would  be  expected,  proportional  to  the 
electromotive  force  impressed  upon  the  electrodes. 

A  large  majority  of  substances  when  brought  in  contact 
with  water  become  charged  negatively  —  the  water  posi- 
tively. Under  these  conditions  water  would  of  course  go  to 
the  cathode.  If  the  substance  hi  question,  when  brought  in 
contact  with  water,  should  be  charged  positively,  the  water 
would  be  charged  negatively  and  go  to  the  anode. 

The  effect  of  electrolytes  —  acids,  bases  and  salts  —  on 
electrical  endosmosis  has  been  studied  by  Perrin  and 
others.  He  found  that  acids  render  negative  diaphragms 
less  negative  and  positive  diaphragms  more  positive.  Alka- 
lies, on  the  contrary,  render  negative  diaphragms  more 
negative,  and  positive  diaphragms  less  positive.  Very 
small  concentrations  of  electrolytes  produce  a  marked 
effect  on  electrical  endosmosis.  A  relation  of  interest  and 
of  some  importance  has  been  worked  out  between  the 
valency  of  the  ions  and  the  action  of  salts  yielding  these 
ions  on  the  difference  in  potential  which  exists  between 
the  liquid  that  is  present  and  the  solid.  When  the  dia- 
phragm is  negative,  the  valence  of  the  cation  present  condi- 
tions the  difference  in  potential.  When  the  diaphragm 
is  positive,  the  valence  of  the  anion  determines  the  poten- 
tial difference. 

These  relations  are  valuable  in  connection  with  the 
precipitation  of  colloids  by  electrolytes. 

Cataphoresis.  —  We  have  seen  that  the  movements  of 

1  Wiedemann:  Elektrizim,  I,  993  (1893). 

«  Pogg.  Ann.,  87,  321  (1852).  4  Ges.  Abhl  I,  855. 

«  Ibid.,  113,  513  (1861).  •  Journ.  Chim.  Phys.,  2,  601  (1904). 


COLLOIDAL  SOLUTIONS  255 

liquids  through  diaphragms  and  capillary  tubes  under 
an  impressed  electromotive  force  is  known  as  electrical 
endosmosis.  The  movements  of  the  solid  suspended 
particles  through  the  liquids  in  which  they  are  suspended, 
under  the  action  of  an  external  electromotive  force,  is 
known  as  cataphoresis.  The  migration  of  the  suspended 
parts  to  the  poles  is,  in  a  sense,  analogous  to  the  migra- 
tion of  the  ions  in  a  true  solution  when  an  electromotive 
force  is  impressed  upon  the  solution. 

This  phenomenon  seems  to  have  been  observed  first  hi 
the  cases  of  arsenic  sulphide  and  ferric  hydroxide  sols  by 
Lander  and  Picton.1 

The  negative  sok  or  those  which  migrate  to  the  positive 
pole,  or  anode,  include  the  metals  silver,  gold,  platinum, 
etc.,  and  most  of  their  compounds  except  their  oxides  and 
hydroxides;  the  trisulphides  of  arsenic  and  antimony; 
the  non-metals  sulphur,  selenium  and  tellurium  and  such 
colloids  as  gamboge,  mastic,  eosin,  aniline  blue,  indigo,  etc. 

The  positive  sols  or  those  which  move  to  the  negative 
pole  or  cathode,  include  the  hydroxides  of  aluminium,  iron, 
chromium,  cadmium,  zinc,  zirconium,  etc.;  colloidal  solu- 
tions of  methyl  violet,  methylene  blue,  etc.,  ^  and  such 
metals  as  iron,  lead  and  bismuth. 

The  determination  of  the  velocities  with  which  the 
particles  move  hi  cataphoresis  is  not  a  difficult  matter. 
The  method  is  analogous  to  that  employed  in  determining 
the  absolute  velocities  with  which  ions  move.  Electrodes 
are  inserted  into  the  two  ends  of  a  vertical  tube  containing 
the  sol,  and  the  current  passed  in  the  direction  so  that  the 
colloid  will  move  to  the  electrode  in  the  lower  end  of  the 
tube.  In  settling,  the  colloid  possesses  a  sharp  upper  sur- 
face, which  can  be  seen  readily  and  with  sufficient  accuracy. 
The  distance  which  it  settles  in  a  given  tune,  under  a  given 
fall  hi  the  potential,  can  easily  be  measured. 

Cotton  and  Mouton2  have  measured  the  velocity  with 

1  Journ.  Chem.  Soc.,  61,  148  (1892). 
1  Journ.  Chim.  Phys.,  4,  365  (1906). 


256  THE  NATURE  OF  SOLUTION 

which  a  single  particle  moves  by  observing  it  under  a 
microscope  provided  with  a  micrometer  eyepiece.  In  this 
way  the  distance  traveled  in  a  given  time  under  a  given 
potential  gradient  could  be  readily  ascertained. 

In  case  the  particle  is  submicroscopic,  the  ultramicro- 
scope  could  and  has  been  used. 

Velocities  with  Which  the  Particles  Move.  —  A  few 
of  the  results1  obtained  in  the  study  of  cataphoresis  are 
given  below. 

Velocity  cm. /sec. 

Diameter  of  for  unit  potential 

Suspended  substance                           the  parts  gradient 

Lycopodium  35  M  25    X  10~* 

Arsenic  trisulphide  50  MM  (about)  22     X  lO"6 

Quartz  1 M  30    X  10~5 

Berlin  blue  (colloid)  <  100  MM  40    X  lO"6 

Gold  (colloid)  <  100  MM  40    X  lO'6 

Platinum  (colloid)  <  100  MM  30    X 10"5 

Silver  (colloid)  <  100  MM  23.6  X  lO"5 

Bismuth  (colloid)  <  100  MM  11.0X1Q-5 

Lead  (colloid)  <  100  MM  12.0  X  lO^5 

Iron  (colloid)  <  100  MM  19.0  X  lO"6 

Ferric  hydroxide  (colloid)  <  100  MM  30.0  X  10~5 

It  is  of  interest  to  note  that  the  velocities  of  the  par- 
ticles are  nearly  constant,  independent  of  their  nature  and 
size. 

We  have  seen  what  is  the  effect  of  the  addition  of  elec- 
trolytes to  electrical  endosmosis.  The  same  question  arises 
here;  but  since  electrolytes  precipitate  so  many  colloidal 
solutions,  it  is  very  difficult,  not  to  say  impossible,  to  an- 
swer it  in  any  broad  way. 

Cataphoresis  with  Emulsoids.  —  The  particles  of  emul- 
soids  under  the  action  of  an  electromotive  force  also  move, 
but  the  motions  are  less  definite  and  the  phenomena  more 
complicated  than  with  suspensoids.  The  emulsoids  are 
more  likely  to  be  precipitated,  and  the  particles  move  more 
slowly  than  with  suspensoids.  The  migration  velocity  of 
the  particles  of  gelatine  is  only  25x10  ~5. 

The  same  emulsoid  particle  under  one  set  of  condi- 
tions may  move  in  one  direction,  under  other  conditions,  in 

1  Freundlich:  Kapittarchemie,  p.  234  (1909). 


COLLOIDAL  SOLUTIONS  257 

the  opposite  direction,  showing  that  it  may  be  charged  either 
positively  or  negatively.  Thus,  if  the  solutions  are  neutral, 
egg-albumin  does  not  move  in  either  direction,  showing  that 
it  is  electrically  neutral.  Perrin1  has  shown  that  in  alkaline 
solution  it  goes  to  the  anode,  having  a  negative  charge;  in 
acid  solution  to  the  cathode,  having  a  positive  charge.  The 
attempt  to  explain  these  phenomena  as  due  to  the  ampho- 
teric  character  of  albumin  meets  with  objections. 

Electrical  Properties  of  Colloids  in  Non-aqueous  Sol- 
vents. —  The  electrical  phenomena  hi  non-aqueous  solvents 
are  in  many  cases  more  complex  than  in  water.  Perrin2 
found  that  there  is  no  electrical  endosmosis  in  such 
solvents  as  carbon  disulphide,  chloroform,  ether,  benzene, 
oil  of  turpentine,  etc.  He  observed  pronounced  endos- 
mosis in  such  solvents  as  water,  methyl  alcohol,  ethyl 
alcohol,  acetone,  nitrobenzene,  etc.  It  will  be  recognized 
at  once  that  the  latter  class  of  solvents  have  high  diekctric 
constants,  and  therefore  great  dissociating  power,  hence 
this  relation  is  of  interest  and  importance.  Those  solvents 
with  high  dielectric  constants  show  pronounced  electrical 
properties. 

Baudouin3  showed  that  substances  which  in  water  are 
positive,  are  also  positive  in  methyl  alcohol  —  a  liquid 
which  also  has  a  high  dielectric  constant;  and  the  converse 
is  also  true. 

Quincke4  proved  that  cataphoretic  phenomena  hi  non- 
aqueous  solvents  hi  general  run  parallel  with  electrical 
endosmosis.  Burton5  demonstrated  that  the  effect  of  tem- 
perature on  electrical  phenomena  in  non-aqueous  solvents 
is  largely  the  effect  of  temperature  on  the  viscosity  of 
the  liquid.  The  explanation  of  the  phenomena  of  electri- 
cal endosmosis  and  cataphoresis  belongs  to  the  future. 
We  are  not  yet  sufficiently  familiar  with  the  facts  to  reach 
any  satisfactory  generalization  which  would  correlate  them. 

1  Compt.  Rend.,  136,  1388  (1903).  4  Pogg.  Ann.,  113,  513  (1861). 

»  Jmirn.  Chim.  Phys.,  2, 601  (1904).  B  Phil.  Mag.  [6],  11,  441  ;(1906). 

«  Compt.  Rend.,  138,  898  (1904). 


258  THE  NATURE  OF  SOLUTION 

Electuostenolysis. — This  term  is  applied  to  those  phe- 
nomena which  manifest  themselves  when  an  electric  cur- 
rent  is  passed  through  fine  cracks  filled  with  a  solution  of 
an  electrolyte.  These  phenomena  have  been  studied  by 
Brown.1  Under  these  conditions  metal  often  separates 
in  the  cracks,  and  the  metal  usually  contains  a  large 
amount  of  hydrogen  gas.  Brown  points  out  that  under  the 
conditions  of  such  experiments  —  strong  current  and  high 
resistance  —  the  solutions  of  the  electrolytes  in  the  cracks 
are  highly  heated.  The  liquids  in  question  decompose, 
liberating  hydrogen  gas  which  reduces  the  salts  of  the 
metals. 

Conductivity  of  Suspensions.  —  The  conductivity  of 
suspensions  is  always  greater  than  that  of  the  water  in 
which  the  parts  are  suspended.  Whitney  and  Blake2  puri- 
fied certain  sols  until  their  conductivity  did  not  further 
change  with  dialysis.  They  obtained  the  following  results. 

Conductivity  Conductivity 

of  the  sol.  of  the  water. 

Gold  sol.  12.7  X  10-«  1.5  X  1CT6 

Platinum  sol.  2.9  X  1(T«  1.1  X  10~« 

Ferric  hydroxide  sol.  35.7  X  10~« 

While  the  conductivities  of  the  colloidal  particles  are 
small,  they  are  much  larger  than  that  of  the  water  in 
which  they  are  suspended.  This  raises  the  question 
whether  the  conductivity  may  not  be  due  to  electrolytes 
held  so  firmly  by  the  colloids  that  they  cannot  be  separated 
by  dialysis;  and  this  question  might  be  difficult  to  answer 
were  it  not  for  the  following  fact.  When  the  current  is 
passed  the  colloid  particles  move  to  the  electrodes.  In 
moving  to  the  electrodes  they  must  carry  electricity  and 
therefore  take  part  in  conduction. 

This  same  question  was  tested  by  Duclaux3  in  the  fol- 
lowing manner.  He  filtered  out  the  suspended  particles 
and  determined  the  conductivity  of  the  filtrate.  It  was 

1  Wied.  Ann.,  42,  450  (1891);  44,  473  (1891). 

2  Journ.  Amer.  Chem.  Soc.,  26,  1339  (1904). 
»  Journ.  Chim.  Phys.,  6,  29  (1907). 


COLLOIDAL  SOLUTIONS  259 

found  to  be  less  than  that  of  the  original  sol,  showing  that 
the  colloidal  particles  had  some  conducting  power. 

The  same  point  was  tested  by  Whitney  and  Blake. 
They  determined  the  conductivity  of  a  gold  sol.  They 
then  separated  a  part  of  the  gold  by  cataphoresis,  and 
redetermined  the  conductivity  of  the  sol,  and  found  that 
it  had  decreased.  This  was  repeated  five  times,  with  the 
result  that  as  more  and  more  of  the  gold  was  precipitated 
from  the  sol,  the  conductivity  became  less  and  less.  This 
is  shown  by  the  following  results.1 

Conductivity 

Gold  sol.  dialyzed  13.2  X  10^ 

After  one  cataphoresis  7.7  X  10~* 

After  two  cataphoreses  4.2  X  10"* 

After  three  cataphoreses  2.7  X  10~« 

After  four  cataphoreses  2.1  X 

After  five  cataphoreses  1.8  X 


Taking  all  of  these  results  together,  it  seems  fairly 
certain  that  suspensions  have  some  conductivity. 

Electrical  Properties  of  Emulsions.  —  While  the  con- 
ductivity of  pure  emulsoid  sols  is  always  small,  yet  it  is 
often  appreciable.  Whitney  and  Blake2  measured  the  con- 
ductivity of  the  sol  of  silicic  acid  and  found  the  value  of 
lOOxlO"6.  For  a  gelatine  sol  they  obtained  the  value 
68xlO~6.  The  cataphoretic  phenomena  presented  by 
these  substances  are  hi  keeping  with  their  conductivities. 

Pauli,3  on  the  other  hand,  freed  albumin  very  carefully 
from  electrolytes  and  studied  its  conductivity.  It  was 
found  to  be  practically  a  non-conductor.  It  showed  prac- 
tically no  cataphoresis.  When  a  tension  of  250  volts  was 
impressed  upon  it,  the  particles  did  not  migrate. 

Emulsoids  with  respect  to  conductivity  fall,  then,  into 
two  classes;  those  that  show  quite  appreciable  conduc- 
tivity, and  those  that  show  little  or  none. 

1  J&urn.  Amer.  Chem.  Soc.,  26,  1346  (1904). 

2  Ibid.,  1374  (1904). 

8  Beitr.  chem.  Physiol.  u.  PathoL,  7,  531  (1906). 


260  THE  NATURE  OF  SOLUTION 

PRECIPITATION  OF  COLLOIDS  BY  ELECTROLYTES 

This  brings  us  to  one  of  the  most  interesting  and 
important  chapters  of  colloid  chemistry;  interesting  be- 
cause it  throws  light  on  the  whole  subject  of  precipitation 
in  chemistry,  and  important  as  telling  us  something  about 
the  nature  of  colloids  themselves.  We  shall  consider  first 
the  action  of  electrolytes  on  suspensoids,  and  then  their 
action  on  emulsoids. 

We  have  seen  that  colloids  prepared  by  any  method  are 
more  or  less  unstable.  The  degree  of  instability,  however, 
varies  greatly  from  colloid  to  colloid,  and  for  the  same 
colloid  varies  with  the  method  employed  in  preparing  it. 
This  is  one  roughly  distinguishing  feature  between  col- 
loidal solutions  and  true  solutions.  The  latter  are  in 
general  perfectly  stable  and  would  persist  indefinitely; 
the  former  are  in  a  state  of  metastability. 

Perhaps  the  most  unstable  of  the  colloids  are  the 
metals,  especially  if  these  are  prepared  by  the  Bredig 
method.  They  degenerate  rapidly,  the  metal  clotting  and 
settling  to  the  bottom.  If  these  systems  are  warmed,  the 
clotting  takes  place  still  more  rapidly  —  and  at  once  if  an 
electrolyte  is  added.  • 

On  the  other  hand  there  are  many  sols,  such  as  ferric 
hydroxide,  which  are  relatively  stable  and  will  persist  for  a 
long  time. 

The  precipitation  of  colloids  by  electrolytes  is  usually 
an  irreversible  phenomenon.  The  colloid  particles  clot 
and  form  a  precipitate  and  this  persists,  not  going  back 
again  into  the  colloid  state.  There  are,  however,  ex- 
ceptions to  this  general  relation.  It  has  been  found  that 
when  gold  from  a  gold  sol  was  deposited  by  cataphoresis, 
it  readily  passed  back  again  into  the  water  reproducing  the 
original,  homogeneous  sol.  A  similar  observation  was  made 
by  Linder  and  Picton1  in  connection  with  a  sol  of  arsenic 
trisulphide.  One  part  was  concentrated  by  cataphoresis 

1  Journ.  Chem.  8oc.t  61,  160  (1892). 


COLLOIDAL  SOLUTIONS  261 

and  the  whole  then  allowed  to  stand.  The  colloid  diffused 
back  into  the  liquid  and  re-established  homogeneity. 

Again,  certain  colloids  which  have  been  precipitated 
can  be  restored  to  the  colloid  condition  simply  by  treating 
the  precipitates  with  certain  chemical  reagents.  Thus, 
Whitney  and  Blake1  found  that  precipitated  gold  could  be 
returned  to  the  colloid  condition  by  simply  treating  it 
with  ammonia.  Under  and  Pic  ton2  found  that  if  colloidal 
ferric  hydroxide  had  been  precipitated  by  a  salt  like 
sodium  chloride,  it  was  only  necessary  to  remove  the  salt 
by  careful  washing  hi  order  to  obtain  the  hydroxide  again 
hi  the  colloidal  condition.  Precipitated  arsenic  trisulphide 
was  restored  to  the  colloidal  condition  simply  by  treatment 
with  hydrogen  sulphide. 

Action  of  Electrolytes  on  Colloidal  Suspensions.  —  The 
action  of  electrolytes,  and  the  non-activity  of  non-electro- 
lytes, on  suspensions  can  be  shown  qualitatively  hi  the 
following  way.  Treat  an  aqueous  solution  of  arsenic 
chloride  with  an  aqueous  solution  of  hydrogen  sulphide, 
arsenic  sulphide  will  be  precipitated.  On  the  other  hand, 
treat  an  aqueous  solution  of  arsenic  trioxide  with  an 
aqueous  solution  of  hydrogen  sulphide  and  no  precipitate 
is  formed;  the  system  simply  acquiring  a  yellowish  brown 
color.  Why  this  difference? 

In  the  first  case  the  reaction  proceeds  thus  — 

2  AsCl3+  3H2S=  As2S3+  3HC1. 
In  the  second  case  thus  — 

ASA*  3H2S=  As2S3+  3H2O. 

What  is  the  difference? 

In  the  first  case  an  electrolyte  —  hydrochloric  acid  —  is 
formed,  which  in  the  presence  of  water  dissociates  into 
ions.  In  the  second  case  no  electrolyte  is  formed;  and  we 
shall  see  that  this  is  often  the  determining  factor  —  the 
presence  or  absence  of  ions  or  charged  parts  determining 

1  Journ.  Amer.  Chem.  Soc.,  26,  1341  (1904). 

2  Journ.  Chem.,  Soc.,  61,  114  (1892);  87,  1924  (1905.) 


262  THE  NATURE  OF  SOLUTION 

whether  we  get  a  precipitate  or  a  colloidal  suspension.  If 
this  is  true,  when  we  add  aqueous  solutions  of  an  acid, 
base,  or  salt,  all  of  which  contain  ions,  to  the  colloid  of 
arsenic  trisulphide  formed,  as  described  above,  from  arsenic 
trioxide,  the  sulphide  of  arsenic  should  be  precipitated. 
Such  is  the  fact.  Conversely,  the  addition  of  non-electro- 
lytes such  as  alcohol,  cane  sugar,  etc.  to  the  arsenic  sul- 
phide sol  should  not  produce  any  precipitation,  and  such 
again  is  the  fact.  This  would  make  it  appear  probable  that 
ions  or  charged  parts  are  the  prime  factors  in  determining 
and  conditioning  the  precipitation  of  colloids.  Why  this 
is  the  case  will  be  seen  later. 

Microscopic  and  Ultramicroscopic  Observation  of  the 
Precipitation  of  Colloids.  —  Linder  and  Picton1  have  de- 
scribed the  precipitation  of  a  suspensoid  by  an  electrolyte, 
as  seen  under  the  microscope.  If  the  solution  of  the  elec- 
trolyte is  introduced  on  one  side  of  the  field,  the  precipita- 
tion or  cloudiness  progresses  across  the  field.  The  solid 
particles  at  first  show  active  Brownian  movement,  which, 
as  the  particles  become  larger,  is  less  and  less  pronounced 
and  finally  entirely  disappears;  the  solid  particles  settling 
down  as  a  precipitate. 

The  phenomena  observed  under  the  ultramicroscope 
are  quite  as  interesting.  Take  a  sol  in  which  the  particles 
are  all  amicrons,  or  in  a  state  of  division  which  is  so  small 
that  they  cannot  be  seen  as  separate  reflecting  discs.  If 
a  very  dilute  solution  of  an  electrolyte  is  added,  the  ami- 
crons become  submicrons  which,  on  further  addition  of 
electrolyte,  become  microns,  and  finally  precipitated  par- 
ticles. The  Brownian  movement,  which  in  the  submi- 
cromic  state  is  very  active,  becomes  less  in  the  micromic, 
and  entirely  ceases  when  the  precipitate  proper  forms. 
The  introduction  of  very  small  amounts  of  electrolytes 
may  not  cause  a  precipitate  to  form.  Enough  must  be 
added  to  cause  the  Brownian  movements  of  the  parts  to 
cease.  This  amount,  however,  is  not  independent  of  con- 

1  Journ.  Chem.  Soc.,  87,  1906  (1905). 


COLLOIDAL  SOLUTIONS  263 

ditions.  The  rapidity  with  which  the  electrolyte  is  added  is 
a  conditioning  factor.  The  more  slowly  the  electrolyte  is 
introduced  the  less  precipitation  of  the  colloid;  therefore, 
in  comparing  the  amounts  of  electrolytes  that  are  neces- 
sary and  sufficient  to  effect  complete  precipitation,  we 
must  define  the  conditions  under  which  the  precipitation 
is  to  be  effected. 

Cause  of  the  Precipitation.- —  Before  taking  up  the  re- 
sults that  have  been  obtained,  let  us  raise  and,  if  possible, 
answer  the  question,  why  do  electrolytes  precipitate  col- 
loidal particles?  This  raises  the  further  question,  why  do 
the  colloidal  particles  remain  in  a  state  of  suspension  and 
not  form  a  precipitate?  And  this  in  turn  raises  the  ques- 
tion still  more  fundamental  for  chemistry,  is  a  precipitate 
or  the  colloidal  state  the  natural  condition  of  a  solid  formed 
as  the  result  of  a  chemical  reaction  between  two  or  more 
substances  in  solution? 

Two  substances  such  as  sulphuric  acid  and  barium 
chloride  react  molecule  for  molecule  in  the  sense  of  the 
following  equation: 

H2SO4+  BaCl2=  2HC1+  BaS04 

The  solid  barium  sulphate  when  first  formed  must  be 
either  in  the  molecular  condition,  or  at  most  only  a  few 
molecules  are  aggregated.  This  is  true  in  general  of  the 
formation  of  solids  from  two  or  more  substances  hi  solu- 
tion. The  clotting  or  precipitation  is  a  secondary  phe- 
nomenon, the  primary  condition  of  the  solid  being  a 
suspension  in  the  liquid  or  the  colloidal  state.  We  are 
so  accustomed  in  qualitative  and  quantitative  analysis  to 
deal  with  precipitates  that  we  are  liable  to  look  upon  them 
as  the  condition  to  be  naturally  expected  when  a  solid  is 
formed;  while  just  the  opposite  is  the  case. 

Why  a  Colloid  is  Unstable.  —  Why  does  the  solid  not 
remain  in  the  colloidal  or  suspended  condition? 

We  have  seen  that  the  particles  in  a  suspensoid  are  hi 
general  charged  with  electricity,  and  all  of  the  particles 


264  THE  NATURE  OF  SOLUTION 

in  any  given  suspensoid  are  charged  with  the  same  kind  of 
electricity.  In  some  suspensoids  all  of  the  particles  are 
charged  positively,  and  in  others  all  of  the  particles  are 
charged  negatively.  In  either  case,  since  all  of  the  par- 
ticles in  any  given  suspensoid  have  charges  of  the  same 
sign,  these  naturally  repel  one  another,  and  this  would 
tend  to  keep  the  particles  apart  and  prevent  the  formation 
of  a  clot  or  precipitate.  Acting  counter  to  this  is  the  sur- 
face-tension between  the  solid  particles  and  the  liquid. 
The  action  of  surface-tension  is  always  to  draw  the  object 
up  into  the  smallest  volume  for  a  given  mass.  This  is  why 
mercury  thrown  upon  a  wooden  table  approaches  the  form 
of  little  spheres  —  the  sphere  having  the  smallest  surface  for 
a  given  volume.  So  with  suspended  particles,  surface-ten- 
sion acts  to  diminish  the  surface  for  a  given  volume,  which 
tends  to  draw  the  small  colloidal  particles  together  into 
larger  particles  which  are  the  units  of  clots  or  precipitates. 

We  have  then  these  two  forces  acting  counter  to  one 
another;  and,  keeping  this  fact  hi  mind,  we  can  under- 
stand the  action  of  electrolytes  hi  causing  precipitation. 

Colloidal  solutions  are  exposed  more  and  more  to  the 
air,  which  contains  a  large  number  of  electrolytes.  These 
are  taken  up  in  larger  or  smaller  quantities  and  effect 
precipitation  of  the  colloidal  particles.  This  accounts  in 
part  for  the  instability  of  colloidal  suspensions. 

How  Electrolytes  Act.  —  Take  a  suspensoid  in  which 
all  of  the  particles  are  charged  positively.  The  water  must 
then  be  charged  negatively.  Add  an  electrolyte,  i.e.,  a 
substance  which  in  the  presence  of  water  dissociates  into 
positively  charged  cations  and  negatively  charged  anions. 
The  positively  charged  colloidal  particles  attract  electro- 
statically the  anions  of  the  electrolyte,  and  the  charges 
on  the  particles  are  thus,  as  it  were,  neutralized.  Surface- 
tension  can  then  draw  these  particles  together  and  form 
a  precipitate. 

It  is  well  known  that  under  such  conditions  some  of  the 
electrolyte  is  usually  carried  down  with  the  precipitate, 


COLLOIDAL  SOLUTIONS  265 

and  cannot  be  washed  out  of  it.  This  is  no  doubt  due,  hi 
part,  to  the  fact  that  the  electrostatic  attraction  between 
the  positive  colloid  and  the  anion  of  the  electrolyte  holds 
the  two  so  firmly  together  that  they  cannot  be  separated 
mechanically. 

If  the  colloidal  particles  are  charged  negatively,  exactly 
the  reverse  of  what  is  described  above  occurs.  The  nega- 
tive particles  attract  the  cations  of  the  electrolyte,  and  the 
remainder  of  the  process  is  as  described. 

This  theory  of  the  precipitation  of  colloids  by  elec- 
trolytes we  owe  primarily  to  Burton.1  Let  us  now  see 
what  experimental  evidence  there  is  bearing  upon  it. 

Precipitation  and  the  Valency  of  the  Ion  with  Opposite 
Charge.  —  If  the  above  conclusion  is  correct,  if  precipita- 
tion is  due  to  the  electrical  neutralization  of  the  charged 
colloidal  particle  by  the  ion  with  the  opposite  charge,  then 
there  ought  to  be  a  simple  relation  between  the  numbers  of 
charges  on  the  precipitating  ions  and  the  amounts  of  these 
ions  required  to  effect  the  precipitation  —  a  relation  between 
the  precipitation  quantities  of  the  ions  and  their  valences. 
Freundlich2  has  carried  out  an  elaborate  series  of  experi- 
ments to  test  this  point.  If  the  sol  is  positive  the 
valence  of  the  anion  carrying  a  negative  charge  is  the  con- 
ditioning factor. 

Precipitation  of  a  Positive  Sol.  —  Take  the  positive 
sol,  ferric  hydroxide,  containing  16.3  miUimols  to  the  liter. 
The  following  are  the  concentrations  of  the  electrolytes 
just  necessary  to  produce  complete  precipitation  of  the  sol 
in  a  given  tune  (two  hours);  all  other  conditions,  such  as 
the  relation  of  the  volume  of  the  electrolyte  to  that  of  the 
sol,  being  kept  constant. 

1  Phil.  Mag.,  12,  472  (1906). 

2  Zeit.  phys.  Chem.,  44,  129  (1903);  Kapillarchemie,  p.  352. 


266                          THE  NATURE  OF  SOLUTION 

Millimols 

Electrolyte  in  liter 

K,  Cl  9.03 

K,  Br  12.5 

K,  I  16.2 

K,  NO*  11.9 

Na,  Cl  9.25 


964 

£t  S04  0.204 

T£,  SO4  0.219 

Mg,  SO«  0.217 

It  thus  requires  about  the  same  amounts  of  the  uni- 
valent  anions  to  effect  the  precipitation,  regardless  of  the 
nature  of  the  cations  with  which  they  were  combined  hi  the 
compound. 

When  we  turn  to  the  bivalent  anions  very  different 
amounts  were  required,  but  practically  the  same  amounts 
regardless  of  the  nature  of  the  cations  combined  with  them 
before  the  substance  was  dissolved  and  dissociated. 

Precipitation  of  a  Negative  Sol.  —  Let  us  turn  now  to 
negative  sols,  such  as  the  trisulphide  of  arsenic.  Treat  this 
with  cations  of  different  valences,  and  see  how  much  of 
each  is  required  to  effect  complete  precipitation  of  the  sol. 
Freundlich  carried  out  these  experiments  with  a  sol  contain- 
ing 7.539  millimols  to  the  liter.  The  number  of  millimols  of 
each  electrolyte  is  given  in  the  table  on  the  opposite  page. 

The  results  with  negative  sols  are  as  satisfactory  as 
with  positive.  About  the  same  quantities  of  univalent 
cations  are  required  to  precipitate  the  negative  sol.  That 
larger  quantities  of  potassium  formate  and  potassium 
acetate  are  required  than  of  potassium  chloride  or  nitrate 
is  probably  due,  in  part,  to  the  fact  that  formic  acid  is 
less  dissociated  than  hydrochloric  or  nitric  acids,  and 
acetic  acid  still  less  than  formic.  The  salts  of  these  weaker 


COLLOIDAL  SOLUTIONS  267 

Millimols 

Electrolyte  in  liter 

K,C1  49.5 

K,  NO,  60.0 

K,  HCOO  86.0 

K,  CH^COO  110.0 

Na,  Cl  51.0 

fi,  Cl  58.4 

Mg,  Cl7  0.717 

Mg,  SO4  0.810 

ct,~Cl2  0.649 
ii  _  __ 

Sr,  C12  .    0.635 

i  __  i    ___ 

Ba,  Cl,  0.691 

Ba    Na)a  0.687 


Zn,CU  0.685 

0.093 
0.095 

acids  are  less  dissociated  than  the  salts  of  the  stronger; 
and  consequently,  greater  concentrations  of  the  solutions 
are  required  to  produce  a  given  number  of  cations.  The 
fact  that  more  of  the  acetate  is  required  than  of  the  for- 
mate is  in  keeping  with  this  view;  acetic  acid  being  weaker 
than  formic,  potassium  acetate  is  less  dissociated  for  a 
given  concentration  than  potassium  formate. 

When  we  turn  to  the  bivalent  cations  we  find  very 
much  smaller  amounts  required  to  effect  complete  pre- 
cipitation than  with  the  univalent  cations.  Here  again 
the  amounts  of  the  different  bivalent  cations  required  to 
effect  the  precipitation  are  of  the  same  order  of  magnitude. 

Very  much  less  of  the  trivalent  cation  aluminium  is 
required  than  of  the  bivalent  cations.  Practically  the 
same  amounts  are  necessary  whether  in  the  salt  they  were 
combined  with  the  chlorine  or  with  the  N03  ion. 


268  THE  NATURE  OF  SOLUTION 

Here  again  the  same  general  relation  comes  out  as  in  the 
positive  sol.  The  amount  of  the  ion  with  opposite  sign 
required  to  effect  complete  precipitation  is  independent 
of  the  nature  of  the  ion  with  the  opposite  sign,  and  depend- 
dent  primarily  on  the  valence  of  the  "active"  ion,  or 
the  number  of  electrical  charges  which  it  carries. 

Take  one  further  example,  the  precipitation  of  the 
negative  platinum  sol  containing  0.7  miUiatoms  to  the  liter. 

Electrolyte  Millimols 

in  liter 

N~a,  Cl  2.5 

K,  Cl  2.2 

Ba,  Cl7  0.058 

(UO2,  NO^a  0.065 

j I I _ 

A12>   (S04)3  0.013 

2 

The  same  relations  manifest  themselves  here  again. 
For  cations  of  the  same  valence  essentially  the  same 
amounts  are  required.  The  amount  of  any  cation  neces- 
sary to  effect  complete  precipitation  is  primarily  a  function 
of  its  valence. 

Precipitation  Not  Proportional  to  Valence.  —  If  we 
examine  the  above  results  quantitatively,  we  will  find  that 
while  less  of  bivalent  ions  than  of  univalent,  and  of  tri- 
valent  than  bivalent,  are  necessary  to  effect  the  precipita- 
tions, the  different  amounts  are  not  proportional  to  their 
valences;  i.e.,  it  does  not  require  half  as  much  bivalent 
as  univalent,  and  one-third  as  much  trivalent  as  uni- 
valent ions  to  effect  complete  precipitation.  The  question 
is,  why  do  these  relations  not  hold  quantitatively? 

One  other  important  factor  must  be  taken  into  account. 
In  order  that  the  cation  or  anion  may  neutralize  the 
charge  on  the  colloidal  particle  it  must  be  adsorbed.  We 
have  already  seen  that  when  colloidal  particles  are  pre- 
cipitated by  electrolytes,  some  of  the  electrolyte  is  carried 


COLLOIDAL  SOLUTIONS  269 

down  in  the  precipitate  and  cannot  be  removed  by  any 
amount  of  washing.  This  electrolyte  is  undoubtedly  ad- 
sorbed by  the  colloid. 

If  the  colloid  containing  adsorbed  electrolyte  is  treated 
with  a  solution  of  another  electrolyte,  the  original  cations 
can  be  replaced  by  those  of  the  second  electrolyte,  which 
are  similarly  adsorbed  by  the  colloid. 

The  different  ions  are  adsorbed  with  different  ease,  and 
this  explains  hi  part  why  the  precipitating  quantities  of 
the  ions  with  different  valences  are  not  proportional  to 
their  valences. 

We  have  thus  far  fixed  our  attention  entirely  on  the 
ion  of  the  precipitating  electrolyte,  which  has  an  electrical 
charge  opposite  to  that  on  the  colloidal  particle.  A 
moment's  thought  will  show  that  we  cannot  ignore  the  ion 
of  the  electrolyte  having  the  same  charge  as  the  colloid. 
This'  often  has  the  effect  of  rendering  the  colloid  more 
stable.  Thus,  the  hydrogen  ion  renders  positive  sols 
more  stable,  while  the  hydroxyl  ion  increases  the  stability 
of  negative  sols.  In  the  cases  of  salts,  the  ion  with  the 
same  charge  as  the  sol  tends  to  render  the  sol  more  stable, 
but  this  effect  is  so  small  in  comparison  with  the  precipi- 
tating action  of  the  ion  of  opposite  sign,  that  the  effect 
of  the  valency  of  the  latter  is  not  completely  marked. 

Taking  all  of  these  factors  into  account,  adsorption, 
protecting  effect  of  the  ion  with  the  same  sign,  etc.,  we  can 
see  why  the  quantities  of  the  ions  of  different  valence  which 
are  necessary  to  effect  complete  precipitation  are  not  pro- 
portional to  their  valences,  and  yet  the  explanation  of 
precipitation  already  considered  may  be  perfectly  valid. 

Precipitation  of  Coarser  Suspensions  by  Electrolytes. — 
The  action  of  electrolytes  on  coarse  suspensions  is  strik- 
ingly analogous  to  their  action  on  the  finer  suspensions. 
Take  a  coarse  suspension  of  sulphur.  It  is  charged 
negatively  and  precipitated  by  sodium,  potassium  and 
ammonium  chlorides,  sulphates,  etc.,  having  a  concentra- 
tion of  from  0.1  to  0.6  molecules  in  a  liter.  Magnesium, 


270  THE  NATURE  OF  SOLUTION 

calcium  and  barium  chlorides  effect  the  precipitation  hi 
concentrations  ranging  from  0.01  to  0.02  molecules  per 
liter.  Here  the  same  relation  manifests  itself  between  the 
valence  of  the  ion  with  charge  opposite  to  that  on  the 
colloid  and  its  precipitating  power,  as  with  the  finer  or 
more  dispersed  suspensions.  This  relation  then  seems  to 
be  general,  regardless  of  the  dispersity  of  the  suspensoid. 

Precipitation  of  Emulsoids  by  Electrolytes.  —  Much 
less  is  known  about  the  action  of  salts  on  emulsoids  than 
on  suspensoids,  and  the  phenomena  seem  to  be  more  com- 
plex and  less  amenable  to  generalization.  The  interesting 
and  important  relations  which  we  have  just  been  consider- 
ing with  suspensoids,  fail  to  manifest  themselves  with 
the  emulsoids.  This  may  be  due  in  part  to  the  fact  that 
the  emulsoids  occupy  a  position  intermediate  between 
suspensoids  and  true  solutions,  probably  owing  to  their 
relatively  high  dispersity. 

The  facts  that  have  been  brought  out  in  connection  with 
the  precipitation  of  emulsoids  by  electrolytes  are  almost 
wholly  empirical,  and  will  therefore  be  very  briefly  discussed. 

The  precipitation  of  suspensoids  is  sometimes  a  re- 
versible, and  sometimes  an  irreversible  process.  Under 
certain  conditions  the  precipitate  will  pass  back  into  the 
colloidal  state;  under  others  it  remains  a  precipitate. 

When  emulsoids  are  precipitated  by  the  addition  of 
salts,  the  process  is  sometimes  reversible,  the  precipitate 
passing  back  into  solution  on  the  addition  of  more  water. 

The  effect  of  salts  on  gelatine  is  in  general  to  cause 
it  to  precipitate  if  the  concentration  of  the  salt  solu- 
tions is  sufficiently  great.  This  process  is  irreversible  and 
thus  differs  from  another  transformation  in  the  case  of 
gelatine  which  is  reversible  and  of  far  greater  interest. 
This  is  the  transformation  of  the  sol  to  the  gel,  and  of  the 
gel  to  the  sol.  Gelation  is  effected  by  simply  allowing 
a  concentrated  solution  of  gelatine  obtained  at  a  higher 
temperature,  to  cool.  The  reverse  process,  or  solation,  is 
effected  simply  by  warming  the  gel. 


COLLOIDAL  SOLUTIONS  271 

The  precipitation  of  albumin  by  some  salts  is  for  a  time 
a  reversible  phenomenon,  by  others  non-reversible  from 
the  beginning.  The  reversible  transformation  becomes 
non-reversible  on  standing.  Small  amounts  of  the  heavy 
metals  are  sufficient  to  precipitate  albumin,  while  concen- 
trated solutions  of  salts  of  the  alkalies  and  alkaline  earths 
are  required. 

No  relations  of  value,  or  even  of  interest,  have  been  dis- 
covered in  connection  with  the  precipitation  of  emulsions 
by  salts.  The  power  of  a  salt  to  precipitate  a  suspension 
is,  as  we  have  seen,  a  function  of  the  valency  of  the  ion 
of  the  salt  having  a  charge  opposite  in  sign  to  that  on  the 
colloid.  No  such  relation  manifests  itself  with  emulsions. 
The  valency  of  the  precipitating  ions  seem  to  have  nothing 
to  do  with  the  phenomenon. 

Precipitating  Action  of  one  Colloid  on  Another. — 
Neisser  and  Friedemann,1  and  Biltz2  found  that  suspensoids 
with  like  electrical  charges  do  not,  when  mixed,  have  any 
appreciable  effect  on  one  another.  Suspensoids  with 
opposite  electrical  charges,  however,  precipitate  one  another. 
This  is  just  what  would  be  expected  from  the  action  of 
electrolytes  on  suspensions.  The  ions  having  a  charge 
opposite  to  that  on  the  colloid  are  the  conditioning  factors 
hi  the  precipitation  of  suspensions  by  electrolytes. 

Thus,  when  the  positive  sols,  aluminium,  iron,  chromium, 
etc.,  hydroxides,  are  added  to  the  negative  sols,  arsenic, 
antimony,  cadmium,  etc.,  sulphides,  both  sols  are  more 
or  less  precipitated;  the  amount  of  each  that  is  precipi- 
tated depending  on  the  relative  amounts  of  the  two  sols 
present.  If  negative  sols  are  added  to  negative,  or  posi- 
tive to  positive,  there  is  no  precipitation. 

When  we  were  studying  the  precipitation  of  colloids 
by  electrolytes,  we  saw  that  the  amount  of  the  precipita- 
tion was  conditioned  not  only  by  the  amount  of  the  elec- 
trolyte added  to  the  colloid,  but  also  by  the  way  in  which 

1  Munch.  Med.  Wochenschrift,  No.  11  (1903). 

2  Ber.  d.  chem.  GeseU,  37,  1095  (1904.) 


272  THE  NATURE  OF  SOLUTION 

it  was  added,  whether  all  at  once,  or  slowly,  a  drop  at  a  time. 
The  same  relation  obtains  for  the  precipitation  of  a  col- 
loid by  one  of  opposite  sign.  If  the  one  sol  is  added 
quickly  to  the  other  there  is  a  different  amount  of  pre- 
cipitation than  if  the  two  sols  are  mixed  a  little  at  a  time. 
The  amount  of  precipitation  of  a  sol  by  one  of  opposite 
sign  is  conditioned  fundamentally  by  the  amount  of  the 
one  that  is  added  to  a  given  amount  of  the  other.  This 
is  shown  by  the  following  data  for  colloidal  antimony 
trisulphide  to  which  colloidal  ferric  hydroxide  was  added. 
To  2  cubic  centimeters  of  the  antimony  trisulphide  sol, 
containing  5.6  milligrams,  the  ferric  hydroxide  sol  was 
added.1 

Milligrams  Result  directly  Result  after 
FejOa                               after  mixing  one  hour 

0.8  Cloudy  Nearly  homogeneous 

3.2  Small  flakes  Unchanged 

4.8  Flakes  Liquid  yellow 

6.4  Complete  precipitation  Complete  precipitation 

8.0  Slow  precipitation  Complete  precipitation 

12.8  Cloudy  Slight  precipitation 

20.8  Cloudy  Homogeneous 

A  small  amount  of  the  ferric  hydroxide  produces  little 
or  no  precipitation.  More  produces  complete  precipita- 
tion, while  a  much  larger  quantity  produces  little  or  none. 

There  is  an  amount  of  one  sol  which  is  just  equivalent 
to  another,  and  which  will  just  precipitate  it.  This  has 
been  found  to  be  the  case  when  the  two  sols  are  present  in 
electrically  equivalent  quantities. 

The  reason  why  large  amounts  of  the  one  sol  do  not 
completely  precipitate  the  other,  is  probably  that  the  par- 
ticles which  are  present  in  large  excess  surround  those 
present  in  smaller  quantity  and  practically  eliminate  them 
from  the  field  of  action. 

Michaelis  and  Pincussohn2  studied  the  action  of  one  sol 
on  another  with  different  sign,  under  the  ultra-microscope. 
They  used  sols  in  which  the  particles  had  different  colors, 

1  Biltz,  Ber.  d.  chem.  GeselL,  37,  1106  (1904). 

2  Biochem.  Zeit.,  2,  251  (1906-1907). 


COLLOIDAL  SOLUTIONS  273 

and  observed  that  the  number  of  each  became  smaller  and 
smaller,  and  therefore  they  must  have  united  with  one 
another. 

Action  of  Emulsoid  on  Emulsoid.  —  We  have  seen 
that  when  emulsions  are  treated  with  electrolytes,  the 
phenomena  are  not  so  simple  or  so  clean  cut  as  when  these 
same  substances  act  on  suspensions.  The  same  holds 
true  when  one  emulsion  is  treated  with  another.  In  the 
first  place,  the  emulsions  are  not  charged  positively  and 
negatively  hi  the  same  clean-cut  manner  as  suspensions. 
Positive  emulsions  will  often  precipitate  negative  emulsions, 
but  the  quantitative  relations  that  obtain  with  suspensions 
are  lacking  here.  Emulsions  are  more  like  true  solutions 
than  suspensions,  are  more  readily  adsorbed,  and  are  more 
capable  of  entering  into  chemical  reaction  with  one  another. 
The  phenomena  here,  as  a  whole,  are  more  complex,  and 
no  general  relations  have  thus  far  been  established. 

Action  of  Suspensions  on  Emulsions.  —  When  the  two 
types  of  colloids  having  opposite  signs  are  mixed  hi  certain 
proportions,  a  precipitation  results.  If  they  are  mixed 
in  other  proportions  they  exert  a  protective  action  on  one 
another.  No  general  relations  have  been  established. 

Protective  Action  of  Colloids.  — Lottermoser  and  von 
Meyer1  noted  that  albumin  protects  silver  sols  from  pre- 
cipitation by  electrolytes  while  Zsigmondy2  studied  the 
action  of  various  sols  on  colloidal  gold.  The  latter  took 
a  certain  amount  of  a  solution  of  colloidal  gold  and  added 
varying  amounts  of  emulsions  and  1  cubic  centimeter  of  a 
two-normal  solution  of  sodium  chloride.  He  determined 
the  amount  of  the  emulsion  which  just  prevented  the 
change  of  the  red  gold  into  blue.  This  expressed  in  mil- 
ligrams he  termed  the  "gold  number."  A  few  of  Zsig- 
mondy's  results  are  given  in  the  table  on  the  following  page. 

1  Journ.  prakt.  Chem.,  66,  242  (1897). 
8  Zeit.  anal  Chem.,  40,  697  (1901). 


274  THE  NATURE  OF  SOLUTION 

Emulsion  added  Gold  number 

Gelatine  0.005-0.01 

Casein  0.01 

Egg  albumin  0.1—0.2 

Gum  arabic  0.15-0.25 

Dextrin  10-20 

Potato  starch  25 

Sodium  stearate  10  (at  60°);  0.001  (at  100°) 

Sodium  oleate  0.4  -  1 

Colloidal  solutions  of  the  sulphides  are  protected  as  well 
as  the  colloidal  metals;  and  silicic  acid  as  well  as  the  or- 
ganic colloids  exerts  a  protecting  influence  on  the  metal 
sols. 

•It  may  be  said  in  general  that  when  a  suspension  is 
added  to  an  emulsion,  the  system  has  the  properties  of  the 
emulsion.  This  has  been  explained  by  Bechhold  as  due 
to  the  emulsion  forming  a  coating  around  the  suspension 
and  giving  its  characteristic  properties  to  the  suspension. 

GELS 

According  to  Freundlich1  a  gel  is  a  two-phase  system 
in  which  the  dispersing  agent  is  a  solid,  and  the  dispersed 
phase  a  liquid.  Gels  are  frequently  formed  when  emul- 
soids  are  concentrated  by  evaporation.  They  consist  of 
thin  walls  of  the  amorphous  solid  surrounding  a  space 
filled  with  liquid.  The  structure  of  gels  has  been  studied 
microscopically  by  Butschli,2  in  the  cases  of  gelatine,  albu- 
min, cellulose,  starch,  etc.  It  is  so  fine  that  the  highest 
magnification  must  be  employed. 

The  study  of  gels  is  extremely  difficult,  and  knowledge 
of  them  has  been  accumulated  but  slowly.  They  are  very 
complex,  do  not  lend  themselves  to  the  ordinary  methods 
of  investigation,  are  very  sensitive  to  the  influences  of  for- 
eign substances,  to  changes  in  temperature,  etc.  Some 
progress  has  been  made,  however,  in  the  study  of  gels; 
especially  hi  the  study  of  the  sol-gel  transformation, 

1  Kapillarchemie,  p.  474. 

2  Verh.  Naturwiss.  Med.  Vereins,  Heidelberg,  N.  F.,  6,  89,  230,  360  (1893— 
1894);  6,287  (1900). 


COLLOIDAL  SOLUTIONS  275 

solvation  and  desolvation  of  gels,  and  the  transformation 
of  the  amorphous  solid  into  a  crystalline  solid. 

Physical  Properties  of  Gels.  —  Barus1  has  shown  that 
gels  are  far  more  compressible  than  solids,  the  compressi- 
bility increasing  with  the  temperature.  When  a  gel  is 
treated  with  water  the  resulting  volume  is  less  than  the 
sum  of  the  volumes  of  the  two.  The  density  is  determined 
by  these  volume  relations. 

It  is  well  known  that  the  viscosities  of  gels  are  a  func- 
tion of  then*  concentrations,  the  viscosity  increasing  rapidly 
with  the  concentration. 

The  diffusion  of  substances  hi  gels  is  not  very  different 
from  diffusion  hi  pure  water,  provided  the  gel  is  not  too 
concentrated.  If  the  gels  are  very  concentrated,  diffusion 
of  salts  through  them  is  very  much  slower  than  through 
pure  water. 

These  facts  may  be  explained  as  follows,  if  we  consider 
the  structure  of  gels  to  be  that  of  a  honeycomb,  or  lattice 
work.  In  a  dilute  gel  the  concentration  of  the  solution 
hi  the  open  spaces  (Hohlraumen)  is  approximately  the 
same  as  that  in  the  outer  liquid;  and  the  gel  walls  offer 
little  resistance  to  the  passage  of  the  dissolved  substance. 
Hence  the  rate  of  diffusion  is  nearly  the  same  as  that  hi 
pure  water.  In  concentrated  gels,  however,  the  gel  walls 
offer  decided  resistance  to  the  passage  of  the  dissolved 
substance;  and  hence,  although  the  diffusion  really  takes 
place  hi  the  gelatine-poor  liquid  in  the  open  spaces  (Hohl- 
raumen), the  rate  of  diffusion  is  considerably  diminished. 

The  expansion  of  the  gel  with  rise  hi  temperature  is  of 
the  same  order  of  magnitude  as  the  expansion  of  the 
liquid  component  of  the  gel. 

It  is  well  known  that  glass  and  many  other  substances, 
when  subjected  to  pressure,  show  double  refraction.  The 
properties  of  glass  show  it  to  be  a  very  viscous  liquid 
cooled  below  its  true  freezing-point.  The  question  arises 
whether  other  viscous  liquids  subjected  to  mechanical 

1  Amer.  Journ,  Science,  6,  285  (1898). 


276  THE  NATURE  OF  SOLUTION 

strain  would  not  show  similar  properties.  Such  has  been 
found  to  be  the  case  with  collodion,  gelatine,  etc. 

Solventation  and  Desolventation  of  Gels.  —  It  is  well 
known  that  gels  have  the  power  of  taking  up  and  of  giving 
up  liquids.  This  process  may  or  may  not  be  analogous 
to  the  hydration  and  dehydration  of  molecules  and  ions 
in  the  presence  of  water.  To  distinguish  it  from  the  latter, 
it  seems  better  not  to  use  the  terms  hydration  and  dehydra- 
tion in  water,  and  solvation  and  desolvation  hi  solvents  in 
general;  but  to  employ  such  distinguishing  words  as  sol- 
ventation  and  desolventation.  This  subject  will  be  dis- 
cussed as  Freundlich1  has  done  under  the  two  heads:  the 
less  elastic  and  the  more  elastic  gels. 

Less  Elastic  Gels.  —  The  study  of  these  systems  we 
owe  almost  entirely  to  Van  Bemmelen.2  He  studied  the 
hydrates  formed  by  such  substances  as  aluminium  and 
iron  oxides,  to  see  whether  they  corresponded  to  definite 
chemical  compounds.  It  was  found  that  in  a  desiccating 
atmosphere  such  systems  lose  water  perfectly  continuously, 
and  take  it  up  from  air  containing  much  water  in  the  same 
manner. 

A  definite  hydrate  in  the  presence  of  its  own  decomposi- 
tion products  has  a  definite  vapor-tension,  and  this  remains 
practically  constant  as  long  as  there  is  any  of  this  hydrate 
present.  Such  is  not  the  case  with  these  gels.  As  the 
water  is  removed,  the  vapor-tension  gradually  becomes 
less  and  less  with  increase  in  concentration,  as  hi  the  case 
of  a  true  solution. 

The  hydration  and  dehydration  of  silicic  acid  has  been 
studied  in  much  detail.  The  dehydration  curve  is  for  the 
most  part  continuous,  but  in  places  is  nearly  parallel  with 
one  of  the  axes,  showing  that  over  this  region  there  is  loss 
of  water  without  any  considerable  change  in  vapor-pres- 
sure. The  opalescent  appearance  in  this  region  suggests  a 

1  Kapillarchemie,  p.  486  (1909). 

2  Ber.  d.  chem.  Gesell,  11,  2228   (1878);    Zeit.  anorg.  Chem.,  6,  466 
(1894);  13,  233  (1897);  18,  14,  98  (1898);  30,  265  (1902). 


COLLOIDAL  SOLUTIONS  277 

change  in  the  dispersity,  which  is  also  indicated  by  a  change 
in  the  color. 

The  hydration  curve  hi  part  follows  the  dehydration 
curve,  but  in  places  differs  markedly  from  it.  The  changes 
hi  dispersity,  opalescence  and  color,  the  reverse  of  those  hi 
dehydration,  manifest  themselves. 

It  has  been  found  that  the  previous  history  of  the  gel, 
or  its  hysteresis,  determines  its  properties. 

More  Elastic  Gels.  —  The  class  of  the  more  elastic 
gels  includes  especially  such  compounds  of  carbon  as  albu- 
min, starch,  agar-agar,  etc.  The  swelling  or  imbibition  of 
these  substances  with  water  are  matters  of  physiological 
importance.  The  amount  of  water  which  can  be  taken  up 
by  gelatine  is  very  large  indeed.  Thus,  Von  Schroder1  has 
found  that  a  plate  of  gelatine  weighing  0.904  grams,  when 
exposed  for  eight  days  to  an  atmosphere  saturated  with 
water-vapor,  will  take  up  0.37  grams  of  water.  When 
exposed  for  twenty  days  under  the  same  conditions,  the 
weight  does  not  further  change.  The  water  thus  taken  up 
can  be  removed  by  placing  the  plate  hi  a  dry  atmosphere. 

If  the  plate  containing  all  the  water  which  it  can 
absorb  from  the  atmosphere,  is  plunged  into  water,  it  will 
absorb  a  far  larger  quantity  of  water.  It  was  found  that 
the  plate  mentioned  above,  weighing  1.27  grams,  would,  at 
room  temperature,  take  up  hi  an  hour  5.63  grams  of  water. 
If  such  a  plate  is  allowed  to  remain  hi  water  for  a  day  or 
so,  it  will  take  up  a  maximum  amount  of  water,  and  this 
will  not  still  further  increase.  There  seems,  under  these 
different  conditions,  to  be  a  kind  of  equilibrium  between 
the  plate  and  the  water-vapor  on  the  one  hand,  and 
liquid  water  on  the  other. 

The  volume  of  the  gel  alone,  hi  taking  up  water,  is  greatly 
increased.  Llideking 2  has  shown,  however,  that  when  a 
gel  has  taken  up  water,  the  volume  is  less  than  the  sum 
of  the  volumes  of  the  dry  gel  plus  the  water. 

The  pressure  which  gels  will  produce  when  they  absorb 

1  Zett.  phys.  Chem.AS,  75  (1903),  *  Wied.  Ann.,  36,  552  (1888). 


278  THE  NATURE  OF  SOLUTION 

water  cannot  be  measured  directly,  since  the  volume  of  the 
gel  which  has  taken  up  water  is  less  than  the  sum  of  the 
volumes  of  the  dry  gel  and  the  water.  Von  Schroder1  used 
unglazed  porcelain  cells  through  which  water  would  pass, 
but  the  gel  could  not.  These  played  the  role  of  the  semi- 
permeable  membranes  in  the  measurements  of  osmotic 
pressure.  The  water  passed  into  the  cell,  was  taken  up 
by  the  gel  and  the  resulting  pressure  was  great  enough  to 
cause  the  cell  to  break.  Roughly  quantitative  measure- 
ments of  these  pressures  were  made  by  Reinke,2  who  worked 
out  also  the  moduli  of  elasticity  of  imbibition  when  dif- 
ferent amounts  of  water  had  been  taken  up. 

Some  idea  can  be  gained  of  the  pressure  produced  by 
imbibition  by  taking  gels  which  have  imbibed  water  and 
drying  them  in  contact  with  thin  plates  of  such  solids  as 
glass.  Cover  a  glass  plate  with  a  thin  layer  of  gelatine 
which  has  taken  up  water.  Place  the  whole  system  in  a 
desiccator  and  dry  the  gel.  The  glass  plate  will  be  strongly 
bent,  and  may  even  be  broken.  This  explains  the  observa- 
tion made  by  Cailletet,3  that  glass  plates  on  which  a  film  of 
gelatine  is  allowed  to  dry,  show  double  refraction.  This  is, 
of  course,  due  to  the  strain  produced  by  the  drying  gel. 

Rate  at  Which  Water  is  Taken  Up.  —  The  rate  at  which 
the  water  is  taken  up  by  gels  was  studied  by  Hofmeister,4 
who  found  that  the  water  was  at  first  taken  up  very 
rapidly,  and  then  more  and  more  slowly.  This  is  shown  by 
the  following  results. 

PLATE  OF  AGAR  0.395  MM.  THICK 

Time  in  minutes  Grams  water  imbibed 

5  2.52 

10  3.02 

15  3.37 

20  3.54 

25  3.73 

30  3.89 

1  Zdt.  phys.  Chem.,  45,  75  (1903). 

2  Hansteins  botan.  Abkandl,  4,  1  (1879). 

3  Compt.  Rend.,  134,  400  (1902). 

4  Archiv.  exp.  Pathol  u.  Pharmakol,  27,  395  (1890). 


COLLOIDAL  SOLUTIONS  279 

Similar  results  were  obtained  by  Reinke.1 

The  Effect  of  Temperature  on  the  Rate  of  Imbibition 
is  a  Matter  of  Importance  in  Living  Things.  —  It  has  been 
studied  by  Dimitrievics,2  and  by  Reinke.3  In  one  case 
investigated,  after  six  hours  at  5°  the  imbibition  was  60 
percent  of  the  maximum;  at  35°  it  was  97  percent  of  the 
maximum.  The  total  amount  of  water  taken  up  or  the 
imbibition  maximum  in  the  two  cases  was,  however, 
essentially  the  same,  the  equilibrium  being  practically 
unchanged  by  temperature.  Nevertheless,  the  equilibrium 
was  reached  at  the  higher  temperature  hi  24  hours,  while 
48  hours  were  required  at  the  lower  temperature. 

Heat  Set  Free  in  Imbibition.  —  The  heat  that  is  liber- 
ated when  imbibition  takes  place  is  a  matter  of  the  very 
greatest  importance  hi  understanding  the  nature  of  this 
process,  and  will  prove  to  be  fundamental  to  the  applica- 
tion of  thermodynamics  to  such  processes  when  enough  is 
known  about  them  to  enable  them  to  be  treated  by  mathe- 
matical procedures.  The  first  to  make  quantitative  meas- 
urements of  the  heat  liberated  in  imbibition  were  Wiedemann 
and  Liideking.4  Some  of  the  best  results  were  obtained 
-by  Rodewald.5  A  few  of  his  data  are  given  below. 

100  grams  dry  starch  Heat  in  calories  evolved 

%  water  per  gram  starch 

0.23  28.11 

2.39  22.60 

6.27  15.17 

11.65  8.43 

15.68  5.21 

19.52  2.91 

Most  of  the  heat  is  liberated  hi  the  imbibition  of  the 
first  amounts  of  water.  This  conclusion  was  confirmed 
by  the  extended  investigations  of  Rodewald  and  Kattein6 
on  a  large  number  of  kinds  of  starch.  This  is  strictly  hi 

Hanst.  bot.  AbhandL,  4,  42  (1879). 

Wissenschaft.  prakt.  Unters.  d.  Gebiet  Pflanzenbaues,  I,  75. 

Hanst.  bot.  AbhandL,  4,  83  (1879). 

Wied.  Ann.,  25,  145  (1885). 

Zeit.  phys.  Chem.,  24, 206  (1897).  6  Ibid.,  33,  586  (1900). 


280  THE  NATURE  OF  SOLUTION 

keeping  with  the  fact  that  high  temperatures  are  necessary 
to  remove  the  last  traces  of  water  from  gels. 

Before  leaving  the  subject  of  imbibition,  the  fact  should 
be  mentioned  that  certain  gels  imbibe  liquids  other  than 
water.  Thus,  caoutchouc  imbibes  carbon  disulphide,  ether, 
chloroform,  etc. 

Imbibiton  in  Aqueous  Solutions.  —  As  Freundlich1 
points  out,  imbibition  hi  solution  is  of  fundamental  im- 
portance, because  the  imbibition  of  colloids  is  funda- 
mental for  living  things,  and  the  imbibition  of  colloids  in 
nature  takes  place  hi  solution. 

When  a  gel  is  brought  into  the  presence  of  a  solution 
of  a  salt,  the  salt  distributes  itself  between  the  gel  and 
the  solvent.  In  the  presence  of  different  salts  imbibition 
by  the  colloid  takes  place  with  very  different  rapidity. 
Thus,  Hofmeister2  showed  that  the  chlorides  of  ammonium, 
sodium  and  potassium,  and  the  nitrate  and  bromide  of 
sodium,  accelerate  the  velocity  of  imbibition;  while  the 
nitrate,  sulphate,  tartrate,  etc.,  of  sodium  and  the  poly- 
hydroxyl  organic  compounds  retard  imbibition. 

Wenzel3  found  that  the  sulphocyanate  ion  greatly  ac- 
celerates the  rate  of  imbibition,  while  Wo.  Ostwald  found 
that  whatever  favors  gelation  retards  imbibition. 

Hofmeister4  has  also  shown  that  adsorption  plays  a 
prominent  r61e  in  imbibition.  Thus,  the  dye-stuff  methyl- 
violet  is  strongly  adsorbed  by  gelatine. 

THEORIES  OF  THE  COLLOIDAL  STATE 

Are  Colloids  Solutions?  —  The  first  question  that  sug- 
gests itself  is,  are  colloids  true  solutions?  To  the  eye, 
many  of  them  appear  to  be  true  solutions.  The  systems 
seem  to  be  homogeneous,  and  the  ordinary  microscope 
often  fails  to  reveal  any  heterogeneity,  as  has  been  stated. 
They  have  some  of  the  properties  of  true  solutions  — 

1  Kapillarchemie,  p.  512  (1909). 

2  Arch.  exp.  PathoL  und  Pharmakol.,  28,  210  (1891). 

8  Quellkraft  der  Rhodanate,  Gera  (1886).  *  loc.  tit. 


COLLOIDAL  SOLUTIONS  281 

osmotic  pressure  causing  diffusion,  lowering  of  freezing- 
point,  lowering  of  vapor-tension;  but  they  have  these 
properties  to  only  a  very  slight  extent.  True  solutions 
show  osmotic  pressures  which  obey  the  laws  of  gas  pres- 
sure. Substances  in  true  solution  lower  the  freezing-points 
and  the  boiling-points  of  solvents  in  terms  of  Raoult's 
laws. 

Certain  investigators  have  attempted  to  use  these 
properties  of  colloids  to  determine  the  molecular  weights 
of  the  colloidal  particles,  just  as  we  use  them  with  true 
solutions  to  determine  the  molecular  weights  of  dissolved 
substances.  Gladstone  and  Hibbert  have  applied  the  freez- 
ing-point and  the  boiling-point  methods  to  the  determina- 
tion of  the  molecular  weights  of  such  colloids  as  caramel, 
gum,  ferric  hydroxide,  etc.  Sabanejew  and  Alexandrow 
determined  by  the  freezing-point  method  the  molecular 
weight  of  albumin  to  be  about  15,000. 

Sabanejew  distinguished  between  colloids  with  smaller 
molecular  weights,  i.e.,  less  than  30,000;  and  those  which 
gave  no  appreciable  lowering  of  the  freezing-point,  and 
which  were  therefore  assumed  to  have  very  great  molecular 
weights,  at  least  greater  than  30,000.  These  include  such 
substances  as  ferric  hydroxide,  starch,  silicic  acid  and  the 
like.  The  large  molecules  or  particles  of  colloids  are, 
of  course,  made  up  of  aggregates  of  smaller  molecules. 
Looking  at  the  solution  theory  of  colloids  in  a  broad  way, 
we  may  say  that  it  is  not  sufficient.  Many  properties  of 
true  solutions  are  either  not  present  in  colloidal  solutions  at 
all,  or  are  present  to  only  a  slight  extent;  and,  conversely, 
colloidal  solutions  have  many  properties  not  found  in 
true  solutions.  The  determinations  of  the  molecular 
weights  of  colloids  referred  to  above,  all  assume  that 
Raoult's  laws  for  lowering  of  freezing-point  and  lowering 
of  vapor-tension  hold  for  colloidal  solutions.  This  is  a 
pure  assumption,  and  a  highly  improbable  one.  There- 
fore, in  assigning  a  number  to  the  molecular  weight  of  any 
colloidal  particle,  we  must  remember  that  it  probably 


282  THE  NATURE  OF  SOLUTION 

bears  no  very  close  relation  to  the  facts,  and  what  that 
relation  is  we  at  present  do  not  know,  and  have  no  means 
of  finding  out. 

The  Adsorption  Theory  of  Colloids.  —  We  have  already 
seen  that  adsorption  plays  a  prominent  r61e  in  colloid 
chemistry.  Certain  colloids  adsorb  certain  dye-stuffs,  and 
when  colloids  are  precipitated  by  electrolytes,  some  of  the 
electrolyte  is  carried  down  with  the  precipitate  and  ad- 
heres so  firmly  to  it  that  it  cannot  be  removed  by  washing. 
This  all  suggests  adsorption  phenomena.  It  is  further 
obvious  that  the  colloidal  particles,  being  so  small,  have 
very  large  surfaces,  and  adsorption  phenomena  are,  as  we 
have  seen,  surface  phenomena.  That  adsorption  must 
be  taken  into  account  in  dealing  with  colloids  is  quite  cer- 
tain, but  this  alone  seems  insufficient  to  explain  the  nature 
of  colloidal  solutions. 

The  Suspension  Theory  of  Colloids.  —  The  prevailing 
theory  of  colloids  today,  is  that  they  are  simply  suspen- 
sions of  finely  divided  parts  in  the  liquids.  This  theory  is 
hi  accord  with  the  microscopic  and  ultramicroscopic  study 
of  colloidal  solutions.  The  coarser  suspensions  can  be  seen 
with  the  microscope,  the  finer  with  the  ultramicroscope.  In 
some  colloidal  solutions  the  particles  are  so  fine  that  they 
cannot  be  seen  even  with  the  ultramicroscope. 

Linder  and  Picton  carried  out  the  following  experiment 
bearing  on  the  suspension  theory  of  colloids.  They  fil- 
tered colloids  under  pressure  through  porous  porcelain. 
Some  were  filtered  out  completely  by  the  porcelain,  while 
others  were  so  fine-grained  that  they  passed  on  through. 
Thus,  the  coarser  suspensions  were  separated  from  the 
finer.  In  terms,  then,  of  the  suspension  theory,  the  col- 
loidal particles  are  simply  in  a  state  of  mechanical  sus- 
pension in  the  liquid;  the  properties  of  the  colloid  being 
primarily  a  function  of  the  size  of  the  colloidal  particles. 
In  the  suspensions  the  particles  are  larger  than  in  the 
hydrosols,  and  the  latter  therefore  resemble  the  true  solu- 
tions more  closely  in  their  properties. 


COLLOIDAL  SOLUTIONS  283 

It  should  be  pointed  out,  however,  that  the  electrical 
properties  of  certain  colloids  cannot  be  satisfactorily  ex- 
plained hi  terms  of  the  suspension  theory  alone.  The 
movements  of  the  colloidal  particles  toward  the  pole 
would  indicate  a  slight  solution  and  ionization  of  the  col- 
loid. It  is  easier  to  explain  the  precipitation  of  one  col- 
loid by  another  having  the  opposite  charge,  on  the  theory 
that  there  is  slight  solution  of  the  colloidal  particles.  It 
would  thus  seem  that  the  suspension  theory  of  colloids 
must  be  supplemented  by  the  solution  theory  to  some 
extent,  and  at  least  hi  some  cases,  hi  order  to  interpret 
the  phenomena  presented  by  colloids. 

A  number  of  modifications  of  the  suspension  theory  of 
colloids  have  been  recently  proposed  by  the  leading  workers 
hi  this  field.  One  supplementary  theory  will  be  referred  to. 

What  is  the  condition  of  the  suspended  colloidal  par- 
ticles? Von  Weimarn1  studied  the  precipitate  formed  by 
bringing  together  solutions  of  certain  concentrations.  He 
found  that  they  were  crystalline.  When  the  solutions  were 
more  dilute  or  more  concentrated,  the  crystals  were 
smaller  and  smaller,  and  finally  could  be  seen  only  with 
the  aid  of  the  ultramicroscope.  It  therefore  seems  prob- 
able, thinks  Von  Weimam,  that  when  very  dilute  solutions 
are  brought  together  and  form  a  colloid,  the  colloidal  par- 
ticles are  also  crystalline. 

Just  as  extrapolation  beyond  the  facts  is  hi  general  a 
dangerous  physical  process,  so  here  it  seems  a  dangerous 
procedure  to  conclude  from  the  crystalline  nature  of  pre- 
cipitates, that  the  much  finer-grained  colloidal  particles 
are  also  crystalline.  They  may  be  crystalline,  or  they  may 
not  be  crystalline;  this  remains  to  be  proved. 

BEAKING  OF  COLLOID  CHEMISTRY  ON  OTHER  BRANCHES 

OF  SCIENCE 

The  bearing  of  colloids  on  living  processes  has  already 
been  referred  to.  This  will  now  be  discussed  in  some  de- 

1  Roll.  Zeit.,  2,  76  (1907). 


284  THE  NATURE  OF  SOLUTION 

tail,  since  colloids  are  coming  more  and  more  to  the  front 
in  dealing  with  biological  processes. 

Bearing  of  Colloids  on  Physiological  Chemistry. — 
As  early  as  1858  the  botanist  Nageli  recognized  that  sub- 
stances were  present  in  plant  cells  in  what  he  called 
"special"  condition.  This  was  before  the  colloidal  con- 
dition of  matter  was  known. 

Light  has  been  thrown  on  the  passage  of  food-stuffs 
through  the  walls  of  plant  cells.  These  walls  are  analo- 
gous to  the  membranes  used  by  Graham  for  effecting 
dialysis.  Through  these  substances,  as  Graham  showed, 
crystalloids  such  as  sugar  can  pass,  but  colloids,  such 
as  starch,  albumin,  etc.,  cannot  pass. 

Pauli1  has  studied  albumin  from  the  standpoint  of  col- 
loidal chemistry;  protoplasm  being  a  very  complex  colloidal 
solution.  Colloidal  chemistry  has  played  a  prominent  part 
in  explaining  the  way  in  which  food  is  supplied  to  living 
things.  One  or  two  well-known  physiological  processes 
will  be  considered  in  some  detail. 

Relations  Between  Organic  and  Inorganic  Ferments. — 
The  method  devised  by  Bredig  for  preparing  colloidal  solu- 
tions of  the  metals  has  already  been  discussed.  It  con- 
sists in  bringing  two  bars  of  the  metal  in  question  near 
together  under  water,  and  passing  an  electric  current  be- 
tween the  bars.  Under  these  conditions  the  metal  is  torn 
off  in  a  finely  divided  state,  and  remains  in  the  water  as  a 
suspension. 

Bredig  and  von  Berneck,2  when  they  described  this 
method,  published  the  results  under  the  heading  "  Inorganic 
Ferments."  Let  us  now  see  why  they  used  this  title. 

Very  minute  traces  of  certain  organic  ferments  decom- 
pose hydrogen  dioxide  at  an  appreciable  rate.  A  colloidal 
suspension  of  platinum  containing  a  gram-atomic  weight 
(195.2  grams)  in  seventy  million  liters  of  water  decomposes 
hydrogen  dioxide  slowly. 

1  Kolloidchemische  Studien  am  Eiweiss,  Dresden  (1908). 
8  Zeit.  phys.  Chem.,  31,  258  (1899). 


COLLOIDAL  SOLUTIONS  285 

This  relation  between  colloidal  platinum  and  the  organic 
ferments  is,  in  itself,  not  so  striking.  There  are  many  other 
substances  which,  in  small  quantities,  decompose  hydrogen 
dioxide.  Superoxides  hi  the  finely  divided  condition  decom- 
pose hydrogen  peroxide. 

One  of  the  most  striking  relations  between  the  colloidal 
metals  and  organic  ferments  is  to  be  found  in  the  action  of 
certain  "poisons "  on  both.  Certain  substances  hi  very 
minute  quantity,  diminish  very  appreciably  the  rate  at 
which  they  decompose  hydrogen  dioxide.  If  we  arrange 
these  poisons  hi  the  order  of  then:  poisonous  property  to 
organic  ferments,  and  then  arrange  them  with  respect  to 
their  poisonous  action  on  colloidal  platinum,  the  two 
arrangements  are  practically  identical. 

It  has  been  further  proved  that  the  action  of  the  organic 
ferments  is  catalytic,  i.e.,  a  little  of  the  ferment  effects  a 
large  amount  of  reaction,  and  the  ferment  does  not  itself 
enter  into  the  reaction. 

Bredig  and  von  Berneck,1  Ikeda2  and  Reinders3  showed 
that  colloidal  solutions  of  gold,  platinum,  iridium,  etc., 
are  true  catalyzers.  We  know  that  catalytic  action  is  fun- 
damentally a  surface  phenomenon;  a  catalyzer  catalyzing  the 
more,  the  larger  its  surface.  We  can  now  understand  how 
it  is  that  such  minute  traces  of  certain  substances  "poison" 
both  the  enzymes  and  the  colloidal  metals.  Surface  ten- 
sion is  greatly  affected  by  small  quantities  of  impurities. 
This  is  well  recognized  hi  all  measurements  of  surface- 
tension.  Minute  traces  of  such  substances  as  hydrocyanic 
acid,  hydrogen  sulphide,  mercuric  chloride,  etc.,  probably 
exercise  their  poisonous  property,  by  way  of  their  effect  on 
the  surface-tension  of  the  catalyzer. 

The  analogies  between  the  organic  ferments  and  the 
colloidal  solutions  of  the  metals  are  thus  many  and  ap- 
parently deep-seated.  This  is  another  important  relation 
between  colloidal  chemistry  and  physiological  chemistry. 

1  Zeit.  phys.  Chem.,  31,  258  (1899). 

*  Ibid.,  37,  1  (1901).  »  Ibid.,  37,  323  (1901). 


286  THE  NATURE  OF  SOLUTION 

Action  of  Toxins  on  Antitoxins.  —  Few  questions  in 
physiological  chemistry  have  come  more  to  the  front  in 
recent  years,  than  the  one  dealing  with  the  reaction  be- 
tween toxins  and  antitoxins.  We  recall  especially  the 
discussion  between  Erlich,  on  the  one  hand,  and  Arrhe- 
nius  and  Madsen  on  the  other.  The  former  would  explain 
the  reaction  purely  on  the  basis  of  structural  chemistry. 
The  latter  would  invoke  the  aid  of  physical  chemistry. 
So  far  as  colloidal  chemistry  plays  a  r61e  hi  these  reac- 
tions, the  subject  must  be  discussed  in  this  place. 

Arrhenius1  studied  the  diffusion  constants  of  certain 
toxins  and  antitoxins  in  gelatine.  From  the  small  values 
found  he  showed  that  both  toxins  and  antitoxins  are 
colloids,  the  antitoxin  showing  the  colloidal  nature  even 
more  pronounced  than  the  toxin. 

Madsen2  and  Biltz3  studied  the  action  of  the  antitoxin 
on  the  toxin  of  diphtheria,  assuming  that  the  phenomenon 
is  primarily  one  of  adsorption.  On  this  assumption  they 
calculated  the  amount  of  antitoxin  to  toxin,  and  then  com- 
pared the  results  of  calculation  with  those  of  experiment. 

Free  toxins  in  Toxins  to  antitoxin 

the  solution  Calculated  Observed 

1.2  180  197 

17.2  240  237 

49.8  270  251 

72.8  280  272 

The  relation  between  the  values  calculated  and  those 
found  experimentally  is  as  close  as  could  be  expected  when 
we  consider  the  nature  of  the  phenomena  in  question. 

Arrhenius4  supposed  that  the  action  of  an  antitoxin 
on  a  toxin  was  a  chemical  reaction  analogous  to  the  action 
of  an  acid  on  a  base.  He  deduced  the  equation  for  the 
equilibrium  of  a  weak  acid  acting  on  a  weak  base,  and 
attempted  to  apply  it  to  the  action  of  certain  antitoxins 
on  toxins.  The  constants  calculated  from  the  Arrhenius 
equation  and  those  found  experimentally  are  surprisingly 

1  Immunochemie,  p.17.  3  Mediz.  Naturwiss.  Arch.,  1,  362  (1907). 

*  Ibid.,  p.  131.  4  Immunochemie,  Chapt.  3. 


COLLOIDAL  SOLUTIONS  287 

close,  which  would  point  to  the  general  correctness  of  the 
Arrhenius  idea. 

There  is,  however,  one  fact  which  seems  to  militate 
against  the  view  of  Arrhenius  —  the  specific  nature  of  the 
reaction  between  toxin  and  antitoxin.  Hober  and  Gor- 
don1 have  shown  that  the  amount  of  a  toxin  which  a 
given  amount  of  its  antitoxin  will  neutralize  depends 
upon  how  rapidly  the  one  is  added  to  the  other.  Take  the 
amount  of  antitoxin  which,  when  added  rapidly,  is  just 
sufficient  to  neutralize  the  toxin,  and  add  it  slowly;  it 
will  not  neutralize  all  of  the  toxin,  which  is  shown  by  the 
fact  that  the  mixture  is  still  strongly  toxic. 

This  suggests  the  reactions  between  colloids  with  which 
we  are  now  familiar,  and  would  seem  to  place  this  reaction 
under  the  category  of  colloids.  It  is  a  little  difficult  to 
interpret  this  fact  hi  terms  of  the  view  that  the  reaction 
between  toxins  and  antitoxins  is  strictly  analogous  to 
the  reaction  between  even  weak  acids  and  weak  bases, 
which,  in  principle,  is  independent  of  the  nature  of  both  the 
acid  and  the  base. 

Both  the  chemical  and  the  adsorption  theories  seem  to 
have  much  hi  their  favor,  but  at  present  we  are  not  suf- 
ficiently conversant  with  the  facts  involved  to  enable  us 
to  decide  between  the  two.  It  may  be  shown  here,  as 
in  so  many  other  cases,  that  both  theories,  which  at  pres- 
ent seem  antagonistic,  contain  a  part  of  the  truth. 

Bearing  of  Colloids  on  Pharmacology  and  Pathology. — 
Colloids  are  coming  more  and  more  to  the  front  in  phar- 
macology.2 This  is  especially  true  where  the  colloids  are 
very  fine-grained  and  therefore  approach  more  nearly  to 
true  solutions.  The  hydrosol  of  silver  has  valuable 
antiseptic  properties,  and  the  hydrosol  of  mercury  has 
proved  to  be  useful  in  dermatology;  and  only  the  begin- 
ning seems  to  have  been  made  hi  the  application  of  col- 
loids to  medicine. 

1  Beitrage  chem.  Physiol.  u.  Pathol,  5,  436  (1904). 

1  Forges,  Kottoidchemie  und  Medicin,  Roll  Zeit.,  6,  301  (1909). 


288  THE  NATURE  OF  SOLUTION 

Some  glimpse  of  the  bearing  of  colloids  on  pathology  is 
given  by  the  work  of  Fischers1  on  oedema.  It  was  hither- 
to supposed  that  this  is  due  to  a  rise  in  blood  pressure,  or 
to  the  walls  of  the  vessels  becoming  less  resistant.  He  has 
shown  that  oedema  results  if  water  is  imbibed  by  certain 
colloids. 

It  is  safe  to  predict  that  here  again  only  a  beginning  has 
been  made  hi  the  application  of  colloids  to,  and  their  bearing 
on,  pathology. 

Bearing  of  Colloids  on  Agricultural  Chemistry.  —  It  has 
been  difficult,  not  to  say  impossible,  to  explain  many  of  the 
processes  which  go  on  in  the  soil  by  physical  and  chemical 
means  alone.  Much  valuable  scientific  work  on  the  chemis- 
try of  soils,  of  fertilizers  and  of  plant  growth  has  recently 
been  done.  Colloidal  chemistry  has  proved  invaluable  in 
interpreting  the  results  that  have  been  obtained. 

There  are  present  in  the  soil  as  colloids  the  hydroxides 
of  silicon,  iron  and  aluminium,  and  especially  humus  which 
plays  such  an  important  role  in  plant  growth.  Humus  ad- 
sorbs phosphoric  acid  which  is  so  essential  to  plant  growth. 
The  bases  are  adsorbed  in  the  order  potassium,  sodium,  cal- 
cium, and  magnesium,  potassium  being  the  most  adsorbed; 
and,  of  the  common  constituents  of  soils,  sodium  and  cal- 
cium the  least.  When  potassium  salts  are  added  to  soils 
already  saturated  with  sodium  or  calcium  salts,  the  place  of 
a  part  of  the  latter  is  taken  by  some  of  the  former,  which 
explains  the  tenacity  with  which  soils  hold  their  potash, 
which  is  so  essential  to  the  growth  of  such  a  large  number 
of  plants.  This  is  another  of  those  facts  usually  taken  as  a 
matter  of  course,  or  of  chance,  upon  which  the  present 
economy  of  nature  so  vitally  rests.  The  adsorption  of  the 
various  constituents  of  fertilizers  and  of  manures  by  the 
soils,  and  the  imbibition  of  water  in  the  soils,  are  probably 
directly  connected  with  the  presence  of  colloids  in  the 
soils. 

1  Das  Odem:  Eine  experimentalle  und  theoretische  Untersuchung  der  Phy- 
siologie  und  Pathologic  der  Wasserbindung  in  Organismen  (1910). 


COLLOIDAL  SOLUTIONS  289 

The  question  arises,  whence  came  these  colloids  hi  the 
soils?  The  answer  is,  largely  from  the  decomposition  of  the 
rocks. 

Bearing  of  Colloids  on  Mineralogy.  —  Corun  has  shown 
that  many  minerals  are  really  gels.  Thus,  opal,  bauxite 
and  many  other  minerals  belong  hi  this  class.  These  gels 
are  produced  by  the  weathering  of  the  rocks. 

Colloidal  chemistry  has  proved  very  useful  hi  explaining 
the  color  of  minerals.  Rock  salt  is  colored  blue  by  the 
presence  hi  it  of  colloidal  sodium.  When  minerals  and 
glass  are  subjected  to  the  radiations  from  radium  they  be- 
come colored,  due  probably  to  the  decomposing  action  of 
the  radium  —  radiations  liberating  colloidal  metal  in  the 
mineral  or  hi  the  glass. 

COLLOIDAL  CHEMISTRY  AND  THE   CHEMICAL 
INDUSTRIES 

We  have  seen  that  the  emulsoids  occupy  a  position  inter- 
mediate between  true  solutions  and  suspensoids.  They  have 
some  of  the  properties  of  true  solutions,  but  have  these 
to  a  much  less  extent  than  true  solutions;  i.e.,  osmotic 
pressure,  lowering  of  freezing-point,  lowering  of  vapor- 
tension.  Their  molecular  weights,  while  large  as  compared 
with  most  truly  dissolved  substances,  are  still  apparently 
much  smaller  than  the  weights  of  the  suspended  particles 
in  true  suspensions. 

There  are  certain  substances  which  seem  to  occupy 
positions  intermediate  between  the  emulsoids  and  true 
solutions,  and  these  have  been  termed  by  Freundlich1  and 
others  semi-colloids.  These  substances  have  very  great 
technical  significance.  They  are  also  of  physiological  im- 
portance. When  albumin  is  saponified  it  yields,  according 
to  Paal,2  certain  acids  which  have  properties  interme- 
diate between  true  solutions  and  emulsoids.  Their  mole- 
cular weights  are  between  700  and  800,  and  they  diffuse 
through  parchment.  They  also  have  some  colloidal  proper- 

i  Kapittarchemie,  436  (1909).          2  Ber.  d.  chem.  GeselL,  35,  2195  (1902). 


290  THE  NATURE  OF  SOLUTION 

ties,  and  are  therefore  classed  as  semi-colloids.  Peptone 
also  belongs  in  this  class.  It  is  an  amphoteric  elec- 
trolyte; its  reactions  show  that  some  of  its  molecules  are 
not  only  dissolved,  but  are  ionized.  Its  molecular  weight 
by  the  freezing-point  method  is  between  300  and  500. 

Action  of  Soaps.  —  Soaps  have  been  studied  by  Krafft 1 
and  co-workers,  and  they  have  found  some  remarkable 
relations  between  properties  and  concentrations.  With 
increase  in  the  concentration  of  the  solution  the  boiling- 
point,  instead  of  being  raised,  is  actually  lowered.  The 
foamy  nature  of  the  solution  rendered  boiling-point  deter- 
minations difficult  and,  in  the  more  concentrated  solutions, 
accurate  determinations  impossible.  Smits2  measured  the 
lowering  of  the  vapor-tension  of  water  at  80°  by  sodium 
palmitate.  A  normal  solution  gave  no  lowering  that 
could  be  measured,  while  a  half-normal  solution  produced 
a  lowering  of  1.3  millimeters  of  mercury. 

This  would  indicate  that  in  the  more  concentrated  solu- 
tion we  have  to  do  almost  entirely  with  colloids,  while  as 
the  solution  is  diluted,  these  pass  over  more  and  more  into 
semi-colloids  and  truly  dissolved  substances.  As  the  dilu- 
tion of  the  solution  increases  the  alkali  salts  of  the  fatty 
acids  are  hydrolyzed  by  the  large  amount  of  water  present 
into  the  free  acids  and  the  free  bases,  which  dissolve  and 
show  the  normal  properties  of  solutions  of  electrolytes. 

Krafft  and  Stern3  have  shown  that  free  organic  acids 
are  present  in  solutions  of  soaps.  When  solutions  of  soaps 
of  palmitic  and  stearic  acids  are  warmed,  the  hydrolysis 
is  increased  and  more  of  the  acids  is  set  free.  These  acids, 
being  comparatively  insoluble  in  water,  exist  as  small 
droplets  which  produce  a  milky  solution.  The  above- 
named  investigators  have  shown  that  the  acids  can  be 
extracted  with  toluene.  While  the  soaps  in  aqueous  solu- 
tion show  colloidal  properties,  in  alcohol  they  form  true 

1  Ber.  d.  chem.  GeselL,  27,  1747  (1894);  28,  2566  (1895);  29,  1328  (1896); 
32,  1584  (1899).  3  Ber.  d.  chem.  Gesell.,  27,  1752  (1894). 

2  Zeit.  phys.  Chem.,  45,  608  (1903). 


COLLOIDAL  SOLUTIONS  291 

solutions.  Krafft1  determined  the  molecular  weights  of  a 
number  of  soaps  in  alcohol,  using  the  boiling-point  method, 
and  found  that  they  are  in  the  simplest  molecular  condition. 

Tanning.  —  The  process  of  tanning  is  also  largely  de- 
pendent upon  colloids.  The  albumin  of  the  skin  takes  up 
water,  the  amount  depending  upon  the  nature  of  the 
electrolyte  present.  In  the  presence  of  certain  acids  enor- 
mous amounts  of  water  are  thus  taken  up. 

The  gels  of  the  skin  undoubtedly  adsorb  sols  from  the 
tanning  liquid,  and  the  quality  of  the  leather  produced 
depends  largely  on  the  nature  of  the  sols  hi  the  liquid. 
The  more  readily  adsorbed  colloids  are  deposited  on  the 
surface  of  the  leather;  the  less  readily  adsorbed  diffusing 
into  the  body  of  the  leather.  By  suitable  choice  of  tan- 
ning liquid,  the  skin  is  thus  tanned  to  different  depths. 

Dyeing.  —  Another  technical  process  hi  which  colloids 
play  a  prominent  r61e,  is  the  art  of  dyeing.  Many  of  the 
organic  dye-stuffs  are  colloids,  and  many  of  the  sub- 
stances dyed  exhibit  the  colloidal  property  of  imbibition. 
Among  the  colloidal  dyes2  are  such  important  substances 
as  Congo  red,  Congo  brown,  benzo-purpurine,  azoblue, 
aniline  blue,  etc. 

On  the  other  hand  many  of  the  dye-stuffs  are  true  solu- 
tions, having  large  osmotic  pressures  and  therefore  showing 
true  diffusion.  By  the  boiling-point  method  they  give 
normal  molecular  weights  and  under  the  ultramicroscope 
appear  perfectly  clear.  To  this  class  belong  such  impor- 
tant dyes  as  alizarine  red,  auramine,  fluorescein,  eosin, 
methylene  blue,  safranine,  rhodamine,  etc. 

Lying  intermediate  between  the  dyes  which  are  very 
colloidal  and  those  which  are  true  solutions,  come  certain 
dyes  which  have  only  some  colloidal  properties.  These 
diffuse  fairly  rapidly,  show  considerable  rise  in  boiling 
point,  but  under  the  ultramicroscope  the  suspended  par- 
ticles can  be  seen.  In  this  class  belong  fuchsine,  methyl 

*  Ber.  d.  chem.  GeseU.,  32,  1595  (1899). 

*  Freundlich:  KapiUarchemie,  p.  564  (1909). 


292  THE  NATURE  OF  SOLUTION 

violet,  Nile  blue,  neutral  red,  etc.  Krafft1  has  shown  that 
many  dyes  which  when  dissolved  hi  water  are  either  true 
colloids  or  show  some  colloidal  properties,  become  true 
solutions  when  dissolved  in  alcohol. 

That  colloids  play  a  prominent  part  in  dyeing  is  shown 
by  the  fact  that  mordants  are  frequently  used  to  render 
the  color  more  permanent.  These  mordants,  such  as 
aluminium  chloride,  acetate,  etc.,  are  either  colloids,  or 
are  hydrolyzed  by  the  water  present  and  form  colloids. 
These  colloids  react  with  the  dyes.  They  may  do  so  by 
having  an  electrical  charge  opposite  to  that  of  the  dye, 
and  both  mordant  and  colloid  dye  being  precipitated  hi 
the  fibers  of  the  fabrics  to  be  dyed.  This  brings  us  to  the 
various  suggestions  that  have  been  made  to  account  for 
the  phenomena  of  dyeing. 

Before  the  significance  of  colloids  was  known,  dyeing 
was  explained  on  purely  chemical  grounds.  The  salt  of  the 
dye  was  decomposed,  the  colored  ion  adhering  to  the  fiber. 
Since  dyeing  often  takes  place  under  conditions  where 
chemical  reactions  are  excluded,  this  view  is  not  sufficient. 

After  solutions  hi  solids  as  solvents  became  clearly 
recognized  as  a  type  of  true  solutions  (see  next  chapter), 
the  bearing  of  them  on  dyeing  was  seen.  Witt2  suggested 
that  in  dyeing  solid  solutions  are  formed,  the  dye  playing  a 
role  somewhat  similar  to  the  ether  in  the  well-known 
Kiister3  experiment  with  solid  solutions.  The  dye  divides 
itself  between  the  liquid  that  is  present  and  the  solid 
material  to  be  dyed. 

Some  evidence  hi  favor  of  the  solid  solution  theory  of 
dyeing  has  been  furnished  by  the  microscopic  examination 
of  the  dyed  materials.  The  dye  does  not  seem  to  be  lim- 
ited to  the  surface  of  the  fibers,  but  is  distributed  through 
the  fiber.  This  would  suggest  solution  of  the  dye-stuff 
in  the  solid  fibers. 

1  Ber.  d.  chem.  GeselL,  32,  1608  (1899). 

2  Farberzeit,  15,  1  (1890). 

8  Zeit.  phys.  Chem.,  13,  445  (1894);  17,  357  (1895). 


COLLOIDAL  SOLUTIONS  293 

After  adsorption  began  to  be  studied  quantitatively 
and  analyzed  mathematically,  it  was  found  that  dyeing 
often  obeyed  the  laws  of  adsorption.  The  substances  to 
be  dyed  often  took  up  the  dye-stuff  in  quantities  which, 
under  given  conditions,  were  essentially  the  same  as  when 
the  dye  was  taken  up  by  ordinary  adsorbents.  In  a  word, 
the  adsorption  equations  held  for  the  taking  up  of  many 
dye-stuffs  by  many  substances  which  were  dyed  by  them. 
Schmidt1  and  Freundlich  and  Losev2  showed  that  the 
amounts  of  the  dye  taken  up  from  baths  containing  the 
dye  at  different  concentrations  could  be  calculated  from 
the  equations  deduced  from  adsorption  phenomena. 

Enough  is  known  about  the  phenomena  of  dyeing  to 
force  us  to  conclude  that  they  are  complex.  There 
seems  to  be  no  question  but  that  colloids  play  a  prominent 
role.  Solid  solutions  under  certain  conditions  may  be 
formed.  All  things  considered,  adsorption  seems  to  be 
one  of  the  most  important  phenomena. 

When  we  have  a  comprehensive  theory  of  dyeing,  it 
will  include  what  is  true  in  all  of  these  suggestions,  as 
each  of  them  probably  contains  an  element  of  truth. 

Other  Technical  Applications  of  Colloids.  —  Caoutchouc 
is  a  colloid,  and  in  obtaining  it  from  the  plant  juice  other 
colloids  come  into  play.  The  vulcanization  of  caoutchouc 
is  dependent  upon  adsorption  processes. 

In  the  preparation  of  porcelain,  colloids  come  prom- 
inently into  play.  Clay  contains  gels,  and  in  the  presence 
of  water  the  properties  of  these  substances  manifest  them- 
selves, and  to  their  presence  is  due  to  a  large  extent  the 
plasticity  of  clay.  Colloids  come  prominently  to  the  front 
in  the  coloring  of  glass.  We  can  now  obtain  glass  of  al- 
most any  color.  We  have  glass  which  is  red,  the  color 
being  due  to  the  presence  of  very  fine  colloidal  particles 
of  gold.  Other  red  glass  owes  its  color  to  small  particles 
of  copper.  Other  colors  are  due  to  cobalt  or  uranium 
compounds. 

1  Zeit.  phys.  Chem.,  15,  56  (1894).  *  Ibid.,  59,  284  (1907). 


294  THE  NATURE  OF  SOLUTION 

There  are  still  a  large  number  of  applications  of  colloids 
to  the  industries  which  might  be  mentioned.  Colloids  play 
a  r61e  in  the  setting  of  cement,  in  the  brewing  of  beer,  but 
it  would  lead  us  beyond  the  scope  and  purpose  of  this  chap- 
ter to  consider  in  detail  any  of  these  matters. 

ADSORPTION 

It  has  long  been  known  that  charcoal  has  the  power 
to  take  up  certain  gases  like  carbon  dioxide  in  large  quan- 
tities. It  can  also  adsorb  certain  constituents  from  solu- 
tions. This  fact  has  been  utilized  to  the  greatest  advantage 
in  such  processes  as  the  purification  of  sugar.  The  color- 
hag  matter  from  the  solutions  of  sugar  is  removed  by  the 
charcoal,  and  a  colorless  product  results. 

Charcoal  is  not  alone  in  possessing  this  adsorbing  prop- 
erty. Many  other  substances  in  a  fine  state  of  division, 
i.e.,  exposing  large  surfaces,  have  this  same  property.  It 
is  largely  upon  adsorption  that  dyeing  depends.  The 
soil  particles  adsorb  plant  food,  and  retain  it  at  their  sur- 
faces. The  term  adsorb  seems  to  have  been  introduced 
into  the  literature  by  du  Bois  Reymond.  It  implies  that 
the  substance  taken  up  is  held  at  the  surface  of  the  adsorb- 
ent, and  is  not  in  a  state  of  chemical  combination  with  it. 
On  account  of  its  importance  from  both  a  scientific  and  a 
technical  standpoint,  much  work  has  been  done  recently 
on  the  nature  of  adsorption;  how  it  is  affected  by  changing 
the  nature  and  quantity  of  the  adsorbent;  how  by 
changes  hi  the  nature  and  quantity  of  the  substance  to  be 
adsorbed;  the  effect  of  temperature  on  adsorption,  etc. 

This  is  the  method  which  we  follow  in  the  case  of  true 
solutions.  What  is  the  effect  on  solubility  of  changing 
the  nature  of  the  solvent?  What  the  effect  of  changing 
the  nature  of  the  substance  to  be  dissolved?  How  is  solu- 
bility affected  by  temperature,  etc. 

Nature  of  the  Adsorbent.  —  Take  first  the  nature  of  the 
adsorbent.  How  does  this  affect  the  adsorption?  It  is 
well  known  that  solubility  in  true  solution  depends  as  much 


COLLOIDAL  SOLUTIONS  295 

on  the  nature  of  the  solvent,  as  of  the  dissolved  substance; 
and  from  the  solubility  of  one  substance  hi  a  given  solvent, 
we  can  draw  no  certain  conclusion  as  to  the  solubility  of 
another  substance  in  that  same  solvent. 

The  adsorbing  power  of  solids  presents  simpler  relations. 
If  a  solid  adsorbent  adsorbs  one  substance  more  than  it 
does  another,  then  a  second  solid  adsorbent  will  adsorb 
the  first  substance  more  than  it  will  the  second,  and  so  on; 
showing  a  relation  here  which  fails  to  manifest  itself  in 
the  case  of  true  solutions. 

In  dealing  with  the  effect  of  the  nature  of  the  substance 
adsorbed  when  adsorption  takes  place  hi  solution,  we  must 
take  into  account  two  variables  —  the  solvent  and  the 
dissolved  substance. 

From  aqueous  solutions  charcoal  adsorbs  acids,  bases  and 
salts  only  slightly;  while  organic  substances,  especially  the 
aromatic  compounds,  are  strongly  adsorbed.  From  solu- 
tions hi  organic  solvents  the  dissolved  substances  are  only 
slightly  adsorbed.  The  following  data  will  illustrate  this 
relation;  the  concentration  of  the  benzoic  acid  hi  the 
several  solvents  is  kept  constant  and  is  equal  to  0.01  mol. 
per  liter. 

Amount  adsorbed 
in  millimols 

Benzoic  acid  in  water  3.27 

"     "  benzene  0.54 

"     "  acetone  0.30 

"     "  ether  0.30 

A  general  relation  has  been  worked  out  between  the 
adsorption  of  substances  dissolved  hi  any  given  solvent 
and  the  adsorption  of  that  solvent  when  dissolved  hi  other 
solvents.  Solvents  from  which  the  dissolved  substances 
are  easily  adsorbed,  are,  when  dissolved  in  other  solvents, 
only  slightly  adsorbed.  The  converse  also  holds:  sub- 
stances which  from  solutions  are  strongly  adsorbed,  when 
used  as  solvents  yield  solutions  from  which  the  dissolved 
substances  are  only  slightly  adsorbed.  Thus,  sulphuric  acid, 
is,  as  we  have  seen,  only  slightly  adsorbed  from  aqueous 


296  THE  NATURE  OF  SOLUTION 

solutions.     Therefore,  substances  dissolved  in  this  solvent 
should  show  strong  adsorption,  and  such  is  the  case. 

The  phenomenon  of  adsorption  is  a  very  rapid  one.  The 
equilibrium  between  the  substance  to  be  adsorbed  and 
the  adsorbent  is  quickly  reached.  The  following  data 
for  succinic  acid  adsorbed  by  carbon  will  confirm  this 
conclusion,  the  amounts  of  acid  adsorbed  being  expressed 
in  milliequivalents  per  gram  carbon. 

Minutes  Amount  adsorbed 
5  0.183 

10  0.325 

30  0.752 

60  1.06 

oo  1.16 

Effect  of  Temperature  on  Adsorption.  —  The  effect  of 
temperature  on  the  ordinary  processes  of  true  solution  is 
very  pronounced.  The  temperature  coefficients  of  solu- 
bility are  usually  large.  Most  substances  dissolve  far 
more  at  higher  than  at  lower  temperatures.  For  some 
compounds,  however,  exactly  the  reverse  is  true.  Whether 
the  temperature  coefficient  of  solubility  is  positive  or 
negative,  it  is  usually  of  very  considerable  magnitude. 

The  influence  of  temperature  on  the  adsorption  of  acetic 
acid  from  an  aqueous  solution  by  charcoal  can  be  seen 
from  the  following  data. 

Temperatures  Adsorption 

0°  1.15 

50.2°  0.728 

93.8°  0.466 

While  there  may  be  no  direct  connection  between 
adsorption  phenomena  and  chemical  reactions,  there  is 
certainly  an  indirect  relation.  Many  gases  are  adsorbed 
in  large  quantities  by  certain  solids.  These  gases  are  then 
present  at  the  surfaces  of  the  solids  at  high  concentrations. 
They  may  be  in  a  different  physical,  or  even  chemical  state, 
than  the  free  gases  themselves.  If  oxygen  is  adsorbed 
by  a  solid,  in  this  condition  it  will  often  effect  oxidations 
which  oxygen  gas  will  not  effect.  The  same  is  true  of 


COLLOIDAL  SOLUTIONS  297 

the  activity  of  adsorbed  hydrogen.  It  may  have  greater 
reducing  power  than  hydrogen  gas.  These  examples  are 
sufficient  to  show  that  while  adsorption  as  such  is  prob- 
ably not  a  chemical  phenomenon,  it  has  an  important 
chemical  bearing,  and  is  therefore  of  interest  and  impor- 
tance to  the  chemist.  One  application  of  adsorption  must 
be  referred  to.  It  has  been  shown  by  Dewar1  and  more 
recently  by  Pfund,2  that  the  power  of  certain  solids  to 
adsorb  gases  can  be  used  very  effectively  in  producing  high 
vacua.  This  is  based  upon  the  fact  that  gases  are  adsorbed 
more  readily,  in  proportion  to  the  amount  of  gas  present, 
when  dilute  or  at  low  pressures,  than  when  more  concen- 
trated or  under  higher  pressure. 

Theories  of  Adsorption.  —  When  gases  and  liquids  are 
adsorbed  by  solids  they  disappear  as  such  and  are  con- 
densed on  the  surfaces  of  the  solids.  This  condensation, 
especially  hi  the  case  of  a  gas,  would  seem  to  lead  to  high 
pressures,  and  it  has  been  assumed  that  at  least  adsorbed 
gases  are  under  high  pressures  hi  or  on  the  solid  adsorb- 
ents. This  same  view  of  condensation  has  been  extended 
from  gases  to  liquids. 

It  has  been  pointed  out  that  while  the  vapor-tension  of 
a  compressed  liquid  is  higher  than  that  of  the  same  liquid 
under  normal  pressure,  temperature  remaining  constant, 
the  vapor-tension  of  an  adsorbed  liquid  is  less  than  that  of 
the  same  liquid  hi  the  free  state.  Since  at  present  we  know 
so  little  of  the  real  nature  of  adsorption,  and  of  the  real 
condition  of  and  hi  adsorbed  liquids,  it  is  impossible  to 
judge  of  the  value  of  this  objection. 

The  above  conception  of  the  condition  of  adsorbed  sub- 
stances has  led  to  the  view  that  adsorption  is  fundamentally 
connected  with  surface-tension.  A  study  of  the  phenomena 
connected  with  surface-tension  and  those  of  adsorption 
has  shown,  however,  that  the  two  often  seem  to  run  counter 
to  one  another. 

Other  suggestions  have  been  made  to  account  for  the 

1  Proceed.  R&y.  Soc.,  74,  126  (1904).  2  Phys.  Zeit.,  13,  870  (1912). 


298  THE  NATURE  OF  SOLUTION 

phenomena  of  adsorption.  Lagergren,1  from  the  Van't 
Hoff-le  Chatelier  principle,  deduced  the  relation  between 
the  effect  of  pressure  on  solubility  and  the  adsorption  of 
the  substance,  on  the  assumption  that  adsorption  is  a  con- 
densation of  the  substance  adsorbed  on  the  surface  of  the 
adsorbent.  Those  substances  which  with  increase  in  pressure 
are  more  soluble  would  be  the  more  adsorbed.  This  pres- 
sure theory  of  adsorption  is  not  entirely  satisfactory. 

Van  Bemmelen,2  Vaubel 3  and  others  have  proposed  a 
chemical  theory  of  adsorption.  They  regard  adsorption 
as  a  loose  chemical  combination  between  the  adsorbent  and 
the  substance  adsorbed.  If  there  is  any  kind  of  chemical 
combination  between  the  adsorbent  and  the  substance 
adsorbed,  it  differs  fundamentally  from  ordinary  chemical 
combination.  Adsorption  does  not  obey  the  laws  of  def- 
inite proportions,  and  there  are  other  serious  objections 
to  this  view. 

It  has  been  suggested  that  in  adsorption  there  is  formed 
a  solid  solution  between  the  adsorbent  and  the  substance 
adsorbed.  Such  a  solid  solution  may,  to  some  extent,  be 
formed  after  the  adsorption  has  taken  place;  but  this  does 
not  account  for  adsorption  itself. 

1  Bihg.  K.  Svensk.  Ak.  Hand.,  24,  2,  Nos.  4  and  5  (1899). 

2  Journ.  prakt.  Chem.,  23,  324,  379   (1881);   Zeit.  anorg.  Chem.,  23,  111 
(1900).  •  IUd.,  74,  232  (1906). 


CHAPTER  XIII 

SOLUTIONS  IN  SOLIDS  AS  SOLVENTS 

WE  have  thus  far  studied,  first,  solutions  in  gases  as  the 
solvents  —  gases,  liquids  and  solids  being  the  dissolved 
substances.  We  next  took  up  in  some  detail  and  with 
some  thoroughness  solutions  in  liquids  as  the  solvents  — 
solutions  of  gases  in  liquids,  of  liquids  hi  liquids  and  of 
solids  hi  liquids.  These  were  all  true  solutions.  Then  we 
studied  colloidal  solutions  and  colloidal  suspensions  in 
liquids  as  the  solvents;  and  now  it  remains  to  study  solu- 
tions hi  solids  as  the  solvents. 

With  reference  to  solutions  of  gases  and  liquids  in  solids 
we  know  very  little.  Solids  like  charcoal  dissolve  gases,  and 
the  amount  of  the  gas  dissolved  is  increased  by  increasing 
the  pressure  to  which  it  is  subjected.  Liquids  are  taken 
up  by  solids,  sometimes  hi  very  considerable  quantity 
and  the  liquid  is  often  held  with  great  tenacity.  Glass 
which  has  been  heated  to  redness  in  a  vacuum  for  a 
month,  will  still  give  off  water.  In  reference  to  solutions 
of  gases  and  liquids  in  solids,  other  than  a  few  superficial 
observations  concerning  them,  we  can  say  that  our  igno- 
rance, even  at  present,  is  very  nearly  complete. 

Solutions  of  Solids  Dissolved  in  Solids,  or  Solid  Solu- 
tions Proper.  —  When  we  come  to  solid  solutions  proper, 
or  solutions  of  solids  dissolved  hi  solids,  we  can  say  we 
really  know  something,  and  our  knowledge  is  of  an  inter- 
esting kind.  The  whole  subject  of  solid  solutions  was 
opened  up  in  a  paper  by  Van't  Hoff,1  which  appeared  hi 
1890,  bearing  the  title  "Solid  Solutions." 

That  mixtures  of  solids  exist  which  have  all  the  super- 
ficial appearances  of  being  true  solutions,  is  well  known. 

1  Zeit.  phys.  Chem.,  6,  322  (1890). 


300  THE  NATURE  OF  SOLUTION 

Says  Van't  Hoff 1  "If  we  regard  a  solid  solution  as  a  solid, 
homogeneous  mixture  of  several  substances,  the  composi- 
tion of  which  can  be  changed  without  destroying  the 
homogeneity,  analogous  to  solutions  in  liquids  as  the  sol- 
vent, it  would  not  be  difficult  to  cite  cases  which  belong 
unconditionally  hi  this  category." 

"  It  would  be  expected  that  the  number  of  cases  of 
solids  dissolving  one  another  would  be  smaller  than  of 
liquids;  just  as  the  number  of  cases  of  liquids  dissolving 
one  another  to  an  unlimited  extent  is  smaller  than  of 
gases;  yet  we  have  many  cases  of  two  solids  each 
dissolving  the  other." 

"  We  may  mention  first  'isomorphous  mixtures/  which 
best  illustrate  the  point  in  question.  The  alums  are  good 
examples.  They,  like  liquids,  are  miscible  with  one  another 
in  all  proportions.  There  are  also  cases  known  which  are 
comparable  with  the  conditions  illustrated  by  water  and 
ether  (limited  miscibility) ;  beryllium  sulphate  and  selenate 
being  examples." 

Van't  Hoff  then  takes  up  the  discussion  of  "mixed 
crystals,"  as  illustrating  solid  solutions;  and  concludes 
this  part  of  the  discussion  with  the  following  paragraph. 

"In  this  connection  we  must  not  omit  an  interesting 
experiment  carried  out  by  Spring,2  which  shows  indirectly 
with  considerable  certainty  the  reciprocal  solubility  of 
solids.  When  an  equimolecular  admixture  of  barium  sul- 
phate and  sodium  carbonate  are  subjected  to  pressure,  he 
observed  a  double  decomposition  amounting  to  20  per 
cent,  which,  on  standing,  increased  to  80  per  cent."  This 
can  be  explained  only  on  the  assumption  of  partial  mis- 
cibility of  these  solids. 

Osmotic  Pressure  in  Mixtures  of  Solids.  —  That  solids 
can  mix  with  solids,  forming  apparently  homogeneous  mix- 
tures, is  obvious  from  what  has  been  said.  The  question 
is,  are  these  homogeneous  solids  true  solutions,  or  not? 

This  can  be  answered  only  by  applying  to  them  the 

1  Zeit.  phys.  Chem.,  6,  323  (1890).  2  Bull.  Soc.  Chim.,  44, 166  (1885). 


SOLUTIONS  IN  SOLIDS  AS  SOLVENTS  301 

tests  of  true  solutions.  What  are  the  properties  of  true 
.solutions?  They  must  have  osmotic  pressure.  The  solute 
must  lower  the  freezing-point  of  the  solvent,  and,  if  the 
boiling-point  of  the  solute  is  considerably  below  that  of 
the  solvent,  must  lower  its  vapor-tension.  Do  these 
homogeneous  mixtures  of  solids  show  all  or  any  of  these 
properties?  These  questions  are  answered  by  Van't  Hoff 
hi  the  paper  referred  to  above. 

Take  first  the  question  of  the  existence  of  osmotic 
pressure  in  solid  solutions.  How  can  this  be  tested? 
Obviously  not  by  any  direct  method  of  measurement,  or 
even  of  demonstration.  But  while  we  cannot  demon- 
strate directly  the  existence  of  osmotic  pressure  in  mixtures 
of  solids,  we  can  prove  it  indirectly. 

Osmotic  pressure,  as  we  have  seen,  is  the  cause  of  all 
diffusion.  If  we  could  prove  the  existence  of  diffusion 
in  mixtures  of  solids,  we  would  prove  the  existence  of 
osmotic  pressure  in  such  mixtures.  This  is  just  what 
Van't  Hoff  has  done. 

He  says,1  "We  need  not  be  surprised  that,  whereas 
diffusion  in  gases  takes  place  very  rapidly  and  hi  liquids 
slowly,  in  solids  it  is  so  greatly  retarded  that  the  phe- 
nomenon long  escaped  observation.  Yet  it  is  now  per- 
fectly clear. "  Van't  Hoff  then  refers  to  examples  of 
diffusion  in  mixtures  of  solids.  Spring  had  shown  that 
when  barium  sulphate  and  sodium  carbonate  are  pressed 
together  and  the  pressure  removed,  the  double  decomposi- 
tion proceeded  after  the  removal  of  the  pressure,  until, 
after  seven  days,  80  percent  of  the  sulphate  had  been 
transformed  into  carbonate.  Diffusion  must  have  taken 
place  in  this  mixture  of  solids  after  the  pressure  was 
removed,  bringing  them  into  more  intimate  contact;  and 
consequently  the  reaction  proceeded  until  about  three- 
fourths  of  the  sulphate  was  transformed  into  carbonate. 
Bars  of  iron  heated  in  carbon  take  up  as  much  as  five 
percent  of  the  carbon.  Colson2  showed  that  the  iron  also 

1  Zeit.  phys.  Chem.,  5, 325  (1890).  2  Compt.  re?id.,  93, 1074  (1881). 


302  THE  NATURE  OF  SOLUTION 

penetrated  the  carbon  for  a  considerable  distance  even 
at  250°.  Calcium  was  also  shown  to  diffuse  into  iron, 
rendering  it  brittle. 

Violle1  heated  porcelain  crucibles  in  graphite  and  noted 
that  the  carbon  penetrated  the  porcelain;  and  Marsden2 
showed  that  the  carbon  had  gone  through  the  porcelain. 

Diffusion  of  solid  through  solid  is  also  of  practical 
importance.  Gore  has  shown  that  zinc  objects  electro- 
plated with  a  thin  layer  of  copper,  acquire  on  standing  the 
zinc  color,  due  to  the  diffusion  of  the  zinc  through  the 
copper  at  ordinary  temperatures. 

Further,  it  was  pointed  out  by  Colson  that  platinum 
surrounded  on  all  sides  by  carbon  which  was  free  from 
silicon,  when  heated  in  a  porcelain  crucible  took  up  silicon 
from  the  crucible  although  the  platinum  did  not  come  in 
contact  with  the  crucible  at  any  point.  This  showed  that 
the  silicon  from  the  crucible  had  passed  through  the  car- 
bon into  the  platinum. 

One  of  the  best  examples  of  the  diffusion  of  metals 
into  metals  is  that  brought  out  by  Roberts-Austen,3 
working  at  the  mint  of  England.  He  took  cylinders  of 
lead  and  polished  their  ends.  He  then  took  discs  of  gold 
and  polished  their  faces.  He  tied  the  gold  discs  on  to  the 
lead  plates  and  placed  these  solids  in  the  vault  at  ordinary 
temperatures.  The  two  metals  were  allowed  to  stand  in 
contact  for  about  four  years,  when  it  was  found  that  the 
discs  had  stuck  to  the  lead.  These  gold  discs  were  removed 
and  sections  of  the  lead  cylinders  cut  and  analyzed.  The 
gold  was  found  to  have  penetrated  into  the  lead  nearly 
a  centimeter,  the  lead  sections  nearest  the  gold  plate 
being  richest  in  gold.  Here  is  an  example  of  a  resistant 
metal  like  gold,  diffusing  into  a  metal  like  lead  at  ordinary 
temperatures,  in  quantities  which  can  be  easily  measured. 

Van't  Hoff4  then  attempts  to  demonstrate  directly  the 
existence  of  osmotic  pressure  in  mixtures  of  solids,  but 

1  Compt.  rend.,  94,  28  (1882).  »  Proceed.  Roy.  Soc.,  67,  101  (1900). 

2  Proc.  Edinb.  Soc.,  10,  712.  «  Zeit.  phys.  Chem.,  5,  327  (1890). 


SOLUTIONS  IN  SOLIDS  AS  SOLVENTS       303 

since  his  method  of  proof  does  not  seem  to  be  free  from 
objection,  it  will  not  be  discussed  here. 

Lowering  of  the  Vapor-Tension  of  Solids  by  Other 
Solids.  —  We  have  seen  that  a  second  fundamental  and 
characteristic  property  of  true  solutions  is  the  power  of 
the  dissolved  substance  to  lower  the  vapor-tension  of  the 
solvent,  provided  the  solute  boils  considerably  higher  than 
the  solvent.  The  question  here  is,  have  we  in  mixtures 
of  solids  any  evidence  that  the  one  solid  lowers  the  vapor- 
tension  of  the  other? 

Van't  Hoff  points  out  that  Hauer1  had  shown  that  such 
is  the  case.  Lead  dithionate  is  a  salt  which  readily 
decrepitates.  This  means  a  solid  salt  of  which  the  water 
has  a  relatively  high  vapor-tension  at  ordinary  temper- 
atures —  so  high  that  it  blows  the  salt  to  pieces  when  it 
is  heated.  When  an  isomorphous  crystal  of  lead  and  cal- 
cium or  strontium  dithionate  was  obtained,  this  was  found 
to  decrepitate  much  less  than  the  pure  lead  salt;  showing 
that  the  presence  of  the  calcium  or  strontium  salt  had 
lowered  the  vapor-tension  of  the  lead  salt. 

We  cannot  hope  to  test  the  above  question  in  any 
general  way,  on  account  of  the  very  small  vapor-tensions 
of  solids,  and  the  small  percentage  lowering  of  this  very 
small  quantity  by  the  presence  of  one  solid  dissolved  hi 
another  solid.  Wherever  we  can  test  the  question,  how- 
ever, we  find  that  in  the  mixed  solids  one  has  lowered  the 
vapor-tension  of  the  other. 

Another  example  illustrating  the  same  point  that  was 
referred  to  above,  is  the  lowering  of  the  vapor-tension  of 
iron  alum  by  forming  with  it  an  isomorphous  mixture  of 
aluminium  alum.  Similarly,  the  vapor-tension  of  copper 
formate  is  lowered  by  the  presence  in  the  solid  crystal  of 
barium  or  strontium  formate.  Van't  Hoff  adds,2  "Let 
it  be  clearly  understood  that  we  are  not  dealing  here  with 
a  lowering  of  vapor-tension  caused  by  admixture  with 

1  Verhandl  Kaiserl.  Konigl.  geol.  Reichsan.,  163  (1877). 
»  Zeit.  phys.  Chem.,  5,  330  (1890). 


304  THE  NATURE  OF  SOLUTION 

substances  of  smaller  vapor-tension;  the  mixtures  show, 
by  their  decrepitation,  a  smaller  tension  than  either  of  the 
constituents." 

Lowering  of  the  Freezing-Point  of  a  Solid  by  Another 
Solid.  —  This  brings  us  to  the  third  necessary  and  suf- 
ficient property  to  have  a  true  solution  —  the  lowering  of 
the  freezing-point  of  the  solvent  by  the  dissolved  sub- 
stance. Have  we  any  evidence  that  one  solid  lowers  the 
freezing-point  of  another  solid  with  which  it  is  mixed? 
There  is  an  abundance  of  evidence. 

The  alloys  illustrate  this.  Two  are  especially  well 
known  — Rose's  and  Wood's  fusible  metals.  The  former 
contains  one  part  lead,  one  part  tin  and  two  parts  bis- 
muth; the  latter  contains  two  parts  lead,  one  part  tin, 
four  parts  of  bismuth  and  one  part  of  cadmium. 

Lead  melts  at          327°  Rose's  metal  melts  at    94° 

Tin  melts  at  232°  Wood's  metal  melts  at  71° 

Bismuth  melts  at    268° 
Cadmium  melts  at  321° 

An  alloy1  containing  15  bismuth,  8  lead,  4  tin,  and  3 
cadmium  melts  at  60°. 

We  do  not  need  any  better  illustration  than  this  of  one 
solid  lowering  the  freezing-point  of  another  solid  with 
which  it  is  mixed. 

We  have  thus  shown  that  the  three  fundamental  prop- 
erties which  characterize  solutions  in  liquids  as  solvents, 
are  present  to  some  extent  in  mixtures  of  solids  with  one 
another,  and  we  are  therefore  justified  in  regarding  such 
mixtures  of  solids  as  true  solutions. 

Van't  Hoff,  in  the  paper  already  referred  to,  worked  out 
a  method  for  determining  the  molecular  weights  of  solids 
in  solid  solutions;  but  it  would  lead  us  too  far  to  discuss 
his  method  in  detail  here,  especially  since  it  cannot  be 
regarded  as  having  led  to  results  of  any  great  significance. 

The  same  may  be  said  of  the  method  worked  out  by 
Ktister2  for  determining  the  molecular  weights  of  substances 

1  Erdmann:  Lehrb.  d.  anorgan.  Chemie,  p.  684  (1906). 
1  Zeit.  phys.  Chem.,  13,  445  (1894);  17,  357  (1895). 


SOLUTIONS  IN  SOLIDS  AS  SOLVENTS        305 

in  solid  solutions.  In  the  opinion  of  the  author,  a  per- 
fectly frank  statement  of  the  situation  with  respect  to 
solid  solutions,  and  the  molecular  weights  of  solids  hi 
such  solutions,  would  be  about  as  follows.  We  do  not 
know  with  even  a  reasonable  degree  of  probability  the 
molecular  weight  of  even  the  simplest  substance  hi  the 
solid  state;  and  we  do  not  know  the  molecular  weight  of 
any  solid  hi  a  solid  solution  with  any  certainty. 

Thus  we  see  that,  while  our  ignorance  of  solid  solutions 
cannot  at  present  be  said  to  be  quite  perfect,  still  it  is  not 
widely  removed  from  it.  Our  knowledge  both  of  pure 
solids  and  of  solid  solutions  is  very  meager.  We  have  just 
scratched  the  surface,  so  to  speak,  of  matter  in  the  solid 
state. 


CHAPTER  XIV 

THE  NEWER  HYDRATE  THEORY 

ABOUT  sixteen  years  ago  the  work  which  led  to  the 
present  hydrate  theory  of  aqueous  solution  was  begun  in 
this  laboratory.  From  a  very  simple  beginning,  which  did 
not  have  for  its  object  the  study  of  the  nature  of  solution 
in  general,  the  work  has  widened  in  a  number  of  direc- 
tions. Much  of  this  work  has  been  carried  out  with  the 
aid  of  grants  generously  awarded  by  the  Carnegie  Insti- 
tution of  Washington;  this  and  the  following  chapter 
are  taken  with  the  permission  of  the  Institution  largely 
from  Publication  of  the  Carnegie  Institution  of  Washington 
No.  210,  Chapter  VII. 

Earlier  Work.  —  In  1899  the  author's  co-operators  Ota1 
and  Knight2  brought  to  light  certain  facts  which  could  not 
be  explained  in  terms  of  any  relation  that  was  then  known. 
They  found  that  certain  double  salts,  such  as  double 
chlorides,  nitrates,  sulphates,  cyanides,  etc.,  produced 
abnormally  great  lowering  of  the  freezing-point  of  water 
when  the  solutions  were  concentrated.  What  was  more 
perplexing  was  the  fact  that  the  molecular  depression  of 
the  freezing-point  increased  with  the  concentration  beyond  a 
certain  definite  concentration. 

Similar  results  were  found  for  a  fairly  large  number  of 
salts  by  Jones  and  Chambers,3  arid  by  Chambers  and 
Frazer  working  with  Jones.4  The  salts  studied  by  these 
workers  were  those  that  are  known  to  be  very  hygroscopic, 
to  have  great  power  of  combining  with  water.  The  question 
arose,  what  did  these  results  mean?  At  that  time,  the 
writer  was  antagonistic  to  any  hydrate  theory,  regarding 

1  Amer.  Chem.,Journ.,  22,5  (1899).  8  Ibid.,  23,  89  (1900). 

1  Ibid.,  22,  110  (1899).  4  Ibid.,  23,  512  (1900). 


THE  NEWER  HYDRATE  THEORY          307 

the  ions  in  solution  as  having  an  existence  not  only  in- 
dependent of  one  another,  but  also  independent  of  the 
molecules  of  the  solvent.  This  seemed  to  be  the  view 
which  was  held  at  that  time  also  by  most  of  those  who 
founded  the  new  school  of  chemistry. 

Yet  it  seemed  impossible  to  interpret  the  results  ob- 
tained hi  terms  of  any  other  assumption  than  that  a  part 
of  the  water  present  was  combined  with  the  dissolved  sub- 
stance, and  was  therefore  removed  from  playing  the  r61e  of 
solvent.  Accordingly,  hi  1900,1  the  suggestion  was  vent- 
ured, for  want  of  any  better,  that  hydration  in  aqueous 
solution  would  explain  these  results.  If  a  part  of  the  water 
present  is  combined  with  the  dissolved  substance,  there 
would  be  less  water  acting  as  solvent;  and  since  freezing- 
point  lowering  is  proportional  to  the  ratio  between  the 
number  of  molecules  of  the  solvent  and  of  the  dissolved 
substance,  the  less  solvent  present  the  greater  the  lowering 
of  its  freezing-point.  It  is  one  thing  to  make  a  suggestion 
which  accounts  for  the  known  facts;  it  is  a  very  different 
matter  to  show  that  this  is  the  only  reasonable  suggestion 
which  will  account  for  them,  to  show  that  the  suggestion 
is  true. 

Aided  by  a  grant  from  the  Carnegie  Institution  of 
Washington,  the  author  started  Dr.  Getman2  on  a  more 
or  less  systematic  study  of  the  whole  problem.  The  ques- 
tion arose,  were  the  results  already  obtained  limited  to  a 
few  compounds,  or  types  of  compounds,  or  was  this  a 
general  phenomenon?  To  answer  this  the  study  of  acids, 
bases,  and  salts  in  concentrated  solutions,  especially  by 
the  freezing-point  and  conductivity  methods  was  taken  up 
and  the  refractivities  of  many  solutions  determined. 

Relation  Between  Lowering  of  the  Freezing-Point  of 
Water  and  Water  of  Crystallization  of  the  Dissolved  Sub- 
stance. —  The  work  of  Getman  included  the  study  of  the 

1  Amer.  Chem.  Journ.,  23,  103  (1900). 

2  IUd.,  27,  433  (1902);  31,  303  (1904);  32,  308  (1904);  Zeit.  phys.  Chem., 
46, 244  (1903);  49,  385  (1904);  Phys.  Rev.,  18,  146  (1904);  Ber.  d.  chem.  GeseU., 
37,  1511  (1904). 


308  THE  NATURE  OF  SOLUTION 

lowering  of  the  freezing-point  of  water  produced  by  con- 
centrated solutions  of  the  chlorides,  bromides,  iodides  and 
nitrates  of  a  large  number  of  metals.  The  relation  between 
lowering  of  freezing-point  and  water  of  crystallization 
can  be  seen  very  well  from  the  curves  for  the  chlorides 
and  nitrates.1 

The  nitrates  of  sodium,  potassium,  and  ammonium, 
which  crystallize  without  water,  produce  the  smallest 
lowering  of  the  freezing-point  of  water.  Then  come  the 
nitrate  of  lithium  with  2  molecules  of  water,  calcium  with 
4,  and  a  large  number  of  nitrates  each  with  6  molecules 
of  crystal  water;  all  give  about  the  same  lowering  of  the 
freezing-point.  Finally,  the  three  nitrates  of  aluminium, 
iron,  and  chromium  with  8  and  9  molecules  of  water,  give 
the  greatest  lowering  of  the  freezing-point  of  water. 

Relations  similar  to  the  above  come  out  for  the  chlorides, 
the  bromides  and  the  iodides.2  The  freezing-point  lower- 
ings  of  water  produced  by  them  are  roughly  proportional 
to  the  amounts  of  water  with  which  the  salts  crystallize. 

If,  on  the  other  hand,  we  compare  the  chlorides  with 
the  bromides,  with  the  iodides,  with  the  nitrates,  similar 
relations  manifest  themselves. 

It  was  found  that  chlorides,  bromides,  iodides  and  ni- 
trates which  have  no  water  of  crystallization,  all  produce 
about  the  same  molecular  lowering  of  the  freezing-point  of 
water,  and  this  is  between  3°  and  4°.  With  salts  that 
crystallize  without  water  there  is  only  a  very  slight  increase 
in  the  molecular  lowering  of  the  freezing-point  with 
increase  in  the  concentration  of  the  solution.  The  salts 
of  lithium,  which  crystallize  with  the  same  amounts  of 
water,  give  approximately  the  same  depressions  of  the 
freezing-point. 

If  we  compare  the  salts  of  the  alkaline  earths  that 
crystallize  with  6  molecules  of  water,  we  find  that  they 
produce  approximately  the  same  lowerings;  the  nitrates 

1  See  Carnegie  Institution  of  Washington,  Publication  No.  60,  p.  24  (1907). 

2  Ibid.,  pp.  20-26  (1907). 


THE  NEWER  HYDRATE  THEORY          309 

of  iron  and  aluminium  with  8  and  9  molecules  of  water 
give  greater  lowerings  than  the  corresponding  halogen  salts 
with  6. 

In  the  first  case  we  have  kept  the  acid  constant  and 
compared  with  one  another  the  salts  of  the  different 
metals  with  the  same  acid.  In  the  second  case  we  have 
kept  the  metal  constant,  and  compared  the  salts  of  a 
given  metal  with  different  acids.  In  both  cases  the  rela- 
tion between  lowering  of  the  freezing-point  of  water  by  the 
dissolved  substance  and  water  of  crystallization  of  the  dis- 
solved substance  manifests  itself. 

Those  salts  that  crystallize  with  the  largest  amounts  of 
water  produce  the  greatest  molecular  lowering  of  the 
freezing-point  of  water.  The  work  was  done  with  con- 
centrated solutions,  and  it  has  already  been  pointed  out 
that  for  such  substances  the  molecular  lowering  of  the 
freezing-point  increases  with  the  concentration  of  the 
solution. 

We  must  now  ask,  what  bearing  has  this  relation  on  the 
question  of  hydration  or  non-hydration  hi  aqueous  solution? 
A  moment's  thought  will  show  that  the  bearing  is  a  very 
direct  one.  If  hydrates  exist  in  aqueous  solution,  those 
substances  which  in  such  solutions  would  form  the  most 
complex  hydrates  would  be  the  substances  that  would 
crystallize  from  aqueous  solutions  with  the  largest  amounts 
of  water.  This  is  the  same  as  to  say  that  those  sub- 
stances which,  in  the  presence  of  a  large  amount  of  water, 
have  the  greatest  power  to  combine  with  water,  would, 
other  things  being  equal,  be  the  ones  to  bring  with  them 
out  of  aqueous  solution  the  largest  amounts  of  water  as 
water  of  crystallization. 

We  could  not,  however,  expect  one  of  these  phenomena 
to  be  strictly  a  linear  function  of  the  other,  since  there  are 
undoubtedly  other  factors,  such  as  shape  of  molecules, 
angles  of  crystals,  etc.,  coming  into  play  in  determining 
the  exact  composition  of  crystals. 

That  a  relation  such  as  was  pointed  out  above  holds 


310  THE  NATURE  OF  SOLUTION 

so  well  and  so  generally  for  such  a  large  number  of  sub- 
stances is  very  significant,  and  pointed  to  the  fact  that 
the  suggestion  of  hydration  in  general  in  aqueous  solution 
contained  more  truth  than  was  imagined  when  it  was  first 
suggested. 

Having  found  a  relation  such  as  the  above,  it  seemed 
desirable  to  look  about  for  others  that  would  bear  directly 
or  indirectly  on  the  problem  in  hand.  Before  taking  up 
these,  another  feature  of  the  work  of  Getman  must  be 
briefly  discussed. 

Approximate  Composition  of  the  Hydrates  Formed 
by  Various  Substances  in  Solution.  —  The  line  of  evidence 
just  discussed  seemed  so  strongly  in  favor  of  the  general 
correctness  of  the  view  that  there  is  combination  between 
the  dissolved  substance  and  some  of  the  water  present, 
that  Jones  and  Getman1  undertook  to  calculate  the  approx- 
imate composition  of  the  hydrates  formed  by  the  different 
substances,  and  by  the  same  substance  at  different  dilu- 
tions. 

The  experimental  work  consisted  in  determining  the 
freezing-point  of  the  solution  and,  consequently,  the  depres- 
sion of  the  freezing-point  of  water  produced  by  4iie  dis- 
solved substance  at  the  concentration  in  question.  From 
the  freezing-point  lowering  the  molecular  lowering  was 
calculated. 

The  dissociation  of  the  solution  was  measured  by  means 
of  the  conductivity  method.  Knowing  the  dissociation, 
the  theoretical  molecular  lowering  was  calculated  on  the 
assumption  that  none  of  the  solvent  was  combined  with 
the  dissolved  substance.  The  ratio  of  the  theoretical 
molecular  lowering  to  the  value  found  experimentally 
gave  the  proportion  of  all  the  water  present  that  was 
uncombined.  The  remainder  of  the  water  was,  of  course, 
combined  with  the  dissolved  substance.  The  total  amount 
of  water  present  in  any  given  solution  could  be  readily 
determined.  It  was  only  necessary  to  take  the  specific 

1  Carnegie  Institution  of  Washington,  Publication  No.  60  (1907). 


THE  NEWER  HYDRATE  THEORY         311 

gravity  of  the  solution  by  weighing  a  known  volume  of  it. 
Knowing  the  specific  gravity  and  the  concentration,  it 
was,  of  course,  quite  simple  to  determine  the  total 
amount  of  water  in,  say,  a  liter  of  the  solution.  The  total 
amount  of  water  in  the  solution  and  the  percentage  of 
combined  water  being  known,  the  total  amount  of  com- 
bined water  was  known.  Knowing  the  amount  of  dissolved 
substance  present  in  a  liter  of  the  solution,  and  knowing 
the  total  amount  of  water  combined  with  it,  it  was  a 
simple  matter  to  calculate  from  the  molecular  weights  of 
the  dissolved  substance  and  the  solvent  how  many  mole- 
cules of  water  were  combined  with  one  molecule  of  the  dis- 
solved substance.  The  results  of  such  a  calculation  are 
only  approximations.  In  the  first  place,  the  conductivity 
method  of  measuring  dissociation  is  not  sufficiently  accu- 
rate for  concentrated  solutions,  and  there  is  no  thoroughly 
reliable  method  known  for  this  purpose.  The  error  here 
is,  however,  in  all  probability  not  very  large.  Another 
source  of  error,  which  is  probably  larger,  results  from 
the  assumption  that  Raoult's  law  holds  for  concentrated 
solutions,  i.e.,  that  for  concentrated  solutions  the  lowering 
of  the  freezing-point  is  proportional  to  the  concentration. 
This  is  not  strictly  the  case,  and  we  do  not  know  at 
present  how  wide  the  deviation  from  Raoult's  law  is  hi 
concentrated  solutions. 

Taking  all  of  these  factors  into  account,  it  still  seems 
highly  probable  that,  by  the  method  outlined  above,  we 
can  arrive  at  a  reasonably  close  approximation  to  the 
amount  of  water  combined  with  a  molecule,  or  the  result- 
ing ions,  of  a  dissolved  substance,  under  given  conditions 
of  concentration.  Whatever  objection  may  be  offered  to 
this  method  of  calculating  the  approximate  composition  of 
the  hydrates  existing  hi  aqueous  solution,  it  should  be 
stated  that,  so  far  as  the  writer  knows,  it  is  the  only  general 
method  thus  far  worked  out  for  throwing  any  light  what- 
ever on  this  important  problem.  Jones  and  Getman  applied 
this  method  of  calculating  the  approximate  composition 


312  THE  NATURE  OF  SOLUTION 

of  hydrates  to  about  100  compounds  —  salts,  acids,  and 
organic  substances  —  and  to  about  1,500  solutions  of  these 
substances.  Their  results  have  been  recorded  in  Publica- 
tion No.  60  of  the  Carnegie  Institution  of  Washington. 

A  few  results  for  strongly  hydrated  salts  will  show  the 
order  of  magnitude  of  the  hydration  for  such  substances. 
In  the  first  column  is  given  the  concentration  in  terms  of 
gram-molecular  normal.  In  the  second,  the  number  of 
molecules  of  water  combined  with  one  molecule  of  the  dis- 
solved substance  at  the  concentration  hi  question. 

Magnesium  chloride  Chromium  chloride 

0.15                     52.1  0.1  66.6 

0.20                      49.7  0.2  49.6 

0.25                     43.3  0.3  42.5 

0.38                      33.2  0.4  39.2 

0.50                      29.4  0.5  39.2 

0.61                     28.3  0.6  35.3 

0.93                     25.8  0.7  35.8 

1.4                       21.4  0.9  31.8 

1.8                       19.0  1.0  30.9 

2.3                       17.4  1.5  26.0 

2.0  22.1 

Salts  of  lithium  form  more  complex  hydrates  than 
those  of  sodium  and  potassium.  This  would  be  expected, 
since  lithium  salts  crystallize  with  water,  while  the  salts 
of  the  other  alkalies  in  general  crystallize  without  water. 

Salts  of  potassium  and  ammonium  generally  crystallize 
without  water,  and  these  compounds,  as  would  be  expected, 
combine  with  relatively  little  water  in  aqueous  solution. 

Many  salts  of  sodium  crystallize  without  water,  and 
these  hydrate  very  slightly.  Other  sodium  salts,  such  as 
the  bromide  and  iodide,  crystallize  with  water  and  show 
considerable  hydrating  power  in  solution. 

Salts  of  calcium  crystallize  with  water  and  all  have,  as 
would  be  expected,  large  hydrating  power.  The  halogen 
salts  crystallize  with  6,  the  nitrate  with  4  molecules  of 
water.  The  nitrate  was  found  to  have  less  hydrating 
power  than  the  chloride  or  bromide. 

The  salts  of  strontium  resemble  those  of  calcium,  both 
hi  the  amounts  of  water  with  which  they  crystallize  and 


THE  NEWER  HYDRATE  THEORY         313 

with  which  they  combine  in  aqueous  solution.  Salts  of 
barium  crystallize  with  less  water  and  show  less  hydration 
than  those  of  calcium  and  strontium. 

The  salts  of  magnesium  have  just  about  the  hydration 
that  would  be  expected  from  their  water  of  crystallization. 
The  same  may  be  said  of  the  salt  of  zinc  that  was  studied. 

Cadmium  is  of  special  interest.  Its  halogen  compounds 
crystallize  with  little  or  no  water,  and  although  cadmium 
belongs  hi  the  same  group  with  metals  of  large  hydrating 
power,  its  halogen  salts  combine  with  only  a  small  amount 
of  water.  The  nitrate  of  cadmium  crystallizes  with  4  mole- 
cules of  water  and,  as  could  be  predicted,  shows  consider- 
able hydrating  power. 

The  chloride  and  nitrate  of  magnesium  show  the  hydra- 
tion that  would  be  expected  from  then-  water  of  crystalliza- 
tion. The  same  may  be  said  of  the  salts  of  nickel,  cobalt, 
and  copper. 

The  chlorides  and  nitrates  of  aluminium,  iron,  and 
chromium  crystallize  with  large  amounts  of  water  and  show 
great  hydrating  power. 

The  strong  mineral  acids  show  some  hydrating  power, 
but  the  complexity  of  the  hydrate  formed  by  these  sub- 
stances seems  to  pass  through  a  maximum.  The  acids 
thus  differ  from  the  salts. 

Some  13  non-electrolytes  were  studied  as  to  their 
hydration,  and  none  of  them  showed  any  appreciable 
hydration.  The  same  applies  to  the  organic  acids  that 
were  studied  in  this  connection. 

The  following  general  relations  were  brought  out  by 
the  work  of  Jones  and  Getman.  The  total  amount  of 
water  held  in  combination  by  the  dissolved  substance 
increases  as  the  concentration  of  the  solution  increases. 
From  what  is  known  of  mass  action,  this  would  be  ex- 
pected. 

The  number  of  molecules  of  water  combined  with  one 
molecule  of  the  dissolved  substance  generally  increases 
from  the  most  concentrated  to  the  most  dilute  solution 


314  THE  NATURE  OF  SOLUTION 

studied.  In  some  cases,  however,  the  number  of  mole- 
cules of  combined  water  seems  to  pass  through  a  maxi- 
mum. These  results,  it  is  believed,  give  us  approximately 
the  amounts  of  combined  water  and  certainly  the  relative 
hydrating  powers  of  the  different  compounds  studied. 

Relation  Between  Water  of  Crystallization  and  Tem- 
perature of  Crystallization.  —  Jones  and  Bassett1  worked 
out  the  approximate  composition  of  the  hydrates  formed 
by  a  large  number  of  substances  and  also  the  following 
relation.  The  hydrates,  as  we  have  seen,  are  very  unstable 
systems.  They  are  readily  broken  down  in  solution  with 
rise  in  temperature.  The  hydrates  which  exist  in  solution 
at  ordinary  temperatures  are  much  more  complex  than 
those  which  the  salts  can  bring  with  them  out  of  solution 
as  water  of  crystallization.  The  hydrates  are  more  stable 
and  more  complex  the  lower  the  temperatures.  We  were, 
however,  surprised,  on  examining  the  literature,  to  find 
the  large  number  of  examples  on  record  of  salts  crystalliz- 
ing with  varying  amounts  of  water,  depending  on  the 
temperature  at  which  the  crystals  were  formed. 

A  few  examples  will  be  given  to  bring  out  the  general 
relation  that  the  number  of  molecules  of  water  of  crys- 
tallization is  larger  the  lower  the  temperature  at  which 
the  salt  crystallizes. 


CaCl2        H20  1       *    XT, 

*                      *          1              A  a    -f-TiQ    4-ATv\T\Ai»o4-Tii«fl 

,    MgCl2    6  H2O  \  Elevated  tempera- 

c\  TT  r\  \       "S  i/fic  tcinpGra/iurc 

2  Sn  r      of    crystallization 
^t  n2vj          •  i  11  

8  H2O  /  tures  above  20 
10H2O..        20° 

6H2O  J 

ia  luwex  aiiu  luwei  . 

12H2O.. 

-  10°  to  -  12° 

MnCl2    2H2O.  .  .. 

20° 

MnS04  3H2O.. 

25°  to  31° 

4H20.... 

15° 

4H20... 

25°  to  31° 

6H20.... 

-21° 

5H20... 

15°  to  20° 

11  H2O 

-21°  to  -37° 

7H20... 

0°  or  below  0° 

12  H2O  '.!'.'. 

-48° 

FeCls   anhydrous.. 

80°  and  above 

FeCl,  3^H20... 

20° 

2H2O  

60°  to  80° 

6H20... 

20°  to  -  16° 

2^  H2O.  .  .  . 

40°  to  60° 

These  examples  suffice  to  show  the  general  nature  of 
the  relation  between  water  of  crystallization  and  the  tem- 

1  Carnegie  Institution  of  Washington,  Publication  No.  60  (1907).    Amer. 
Chem.  Journ.,  33,  534  (1905);  34, 291  (1905).  Zeit.  phys.  Chem.,  62, 231  (1905). 


THE  NEWER  HYDRATE  THEORY         315 

perature  at  which  the  salt  crystallizes.  This  relation 
could  have  been  foreseen  as  a  necessary  consequence 
of  the  theory  of  hydrates  in  aqueous  solution,  and  the 
instability  of  those  hydrates  at  higher  temperatures. 

Dissociation  as  Measured  by  the  Freezing-Point 
Method  and  by  the  Conductivity  Method.  —  When  the 
theory  of  electrolytic  dissociation  was  proposed,  it  be- 
came a  problem  to  measure  accurately  the  magnitude  of 
dissociation.  Arrhenius  pointed  out,  in  his  original  epoch- 
making  paper,  that  dissociation  could  be  measured  either 
by  the  freezing-point  or  by  the  conductivity  method. 

Ostwald  so  improved  the  freezing-point  method  that  it 
could  be  used  to  measure  dissociation. 

A  comparison  of  the  data  from  the  freezing-point 
method  with  those  from  the  conductivity  method  showed 
that  dissociation  as  measured  by  the  former  was  slightly 
higher  than  by  the  latter.  (See  p.  131).  The  meaning 
of  this  discrepancy  was  at  that  time  not  understood. 

After  it  had  been  established,  with  reasonable  cer- 
tainty, that  hydration  takes  place  in  aqueous  solution,  a 
possible  explanation  of  this  apparent  discrepancy  pre- 
sented itself.  But  before  offering  this  explanation  it 
seemed  desirable  to  do  more  experimental  work,  having 
this  point  especially  in  mind.  Pearce1  carried  out  in  this 
laboratory  a  very  careful  piece  of  work,  in  which  dissocia- 
tion was  measured  by  the  freezing-point  method  and  also 
by  the  conductivity  method,  and  the  two  sets  of  results 
were  compared.  The  chlorides  of  calcium,  strontium,  mag- 
nesium, barium,  cobalt,  copper,  and  aluminium  and  the 
nitrates  of  calcium,  magnesium,  barium,  cobalt,  nickel  and 
copper  were  studied  as  well  as  sodium  bromide  and  hy- 
drochloric, nitric  and  sulphuric  acids. 

That  hydration  can  explain  the  fact  that  dissociation 
as  measured  by  freezing-point  is  higher  than  as  measured 
by  conductivity  can  be  seen  from  the  following.  The 
combined  water  is  removed  from  the  field  of  action  as 

1  Carnegie  Institution  of  Washington,  Publication  No.  ISO,  p.  57. 


316  THE  NATURE  OF  SOLUTION 

solvent;  only  the  uncombined  water  is  acting  as  solvent. 
Freezing-point  lowering  is  proportional  to  the  ratio  between 
the  number  of  molecules  of  the  dissolved  substance  and  of 
the  solvent.  If  one-fourth  of  the  water  present  is  com- 
bined with  the  dissolved  substance,  the  freezing-point  lower- 
ing would  be  one-third  greater  than  if  all  the  water  were 
present  as  free  water  and  therefore  acting  as  solvent  water. 
Freezing-point  lowering  would  thus  be  affected  proportion- 
ally by  hydration.  Dissociation  of  concentrated  solutions 
calculated  from  the  freezing-point  lowering  would  therefore 
be  much  too  high. 

The  conductivity  of  a  solution  depends  upon  the  num- 
ber of  ions  present  and  their  velocities.  The  number  of 
ions  would  probably  not  be  affected  greatly  by  the  hydra- 
tion, but  their  velocities  would  be.  The  hydrated  ions 
would,  of  course,  move  more  slowly  than  the  unhydrated. 

The  effect  of  hydration  would  obviously  be  more  pro- 
nounced on  freezing-point  lowering,  which  is  proportional 
to  the  amount  of  solvent  present,  than  on  conductivity. 

Temperature  Coefficients  of  Conductivity  and  Hydra- 
tion. —  A  fairly  elaborate  investigation  on  the  con- 
ductivities, dissociation,  and  temperature  coefficients  of 
conductivity  and  dissociation  of  aqueous  solutions  was 
begun  in  the  author's  laboratory  about  15  years  ago. 
The  work,  as  a  whole,  has  been  recently  published  by  the 
Carnegie  Institution  of  Washington.1  The  monograph  in 
question  contains  the  investigations  of  Clover,2  Hosford,3 
Howard,4  Jacobson,5  Kreider,6  Shaeffer,7  Smith,8  Springer,9 
West,10  Wight,11  Wightman,12  and  Winston.13  The  results 
published  in  this  monograph  are  for  about  110  salts, 

Carnegie  Institution  of  Washington,  Publication  No.  170  (1912). 

Amer.  Chem.  Journ.,  43,  187  (1910). 

Ibid.,  46,  240  (1911).  7  Ibid.,  49,  207  (1913). 

Ibid.,  48,  500  (1912).  8  Ibid.,  50,  1  (1913). 

Ibid.,  40,  355  (1908).  »  Ibid.,  48,  411  (1912). 

Ibid.,  46,  282  (1911).  ™  Ibid.,  44,  508  (1910). 
11  Ibid.,  42,  520  (1909);  44,  159  (1910). 

»  Ibid.,  46,  56  (1911);  48,  320  (1912).  «  Ibid.,  46,  368  (1911). 


THE  NEWER  HYDRATE  THEORY         317 

which  were  studied  from  zero  to  65°,  and  from  the  most 
concentrated  solution  that  could  be  used  to  the  dilution, 
hi  most  cases,  of  complete  dissociation.  The  temperature 
coefficients  of  conductivity  were  calculated  both  in  con- 
ductivity units  and  in  percentage. 

Similar  data  were  obtained  for  about  90  of  the  more 
common  organic  acids,  and  the  constants  for  the  weaker 
acids  were  calculated  from  the  Ostwald  dilution  law.  The 
dissociations  of  the  salts  and  acids  at  the  different  tem- 
peratures were  also  calculated  whenever  possible. 

The  temperature  coefficients  of  conductivity  were  calcu- 
lated both  hi  percentage  and  hi  conductivity  units.  A 
study  of  the  temperature  coefficients  of  conductivity, 
expressed  in  conductivity  units,  brought  out  a  relation 
which  had  a  very  direct  bearing  on  the  question  of  hydra- 
tion  hi  aqueous  solution.  This  is  so  important  that  it 
will  be  discussed  here  hi  some  detail. 

The  conductivity  of  a  solution  is  conditioned  by  the 
number  of  ions  present  and  the  velocities  with  which  they 
move.  Rise  hi  temperature  not  only  does  not  increase 
the  number  of  ions  present,  but,  as  is  well  known,  diminishes 
dissociation.  The  effect  of  rise  hi  temperature  increasing 
the  conductivity  of  solutions  is,  then,  due  to  an  increase 
hi  the  velocities  with  which  the  ions  move.  If  the  ion 
is  driven  by  a  constant  force,  its  velocity  would  be  deter- 
mined chiefly  by  the  viscosity  of  the  solvent  and  by  the 
mass  and  size  of  the  ion.  With  rise  hi  temperature  the 
driving  force  would  be  increased.  Rise  hi  temperature 
would  also  decrease  the  viscosity  of  the  solvent.  The  effect 
of  rise  in  temperature  on  both  of  these  factors  would  be  to 
increase  the  velocities  of  the  ions  and,  consequently,  the 
conductivity. 

Another  factor  must,  however,  be  taken  into  account. 
That  many  ions  in  aqueous  solutions  are  strongly  hy- 
drated  seems  now  quite  generally  accepted.  We  have  seen 
that  these  hydrates  are  relatively  unstable  and  break  down 
with  rise  in  temperature.  The  simpler  the  hydrate  formed 


318 


THE  NATURE  OF  SOLUTION 


by  an  ion,  the  smaller  the  mass  of  the  ion;  the  smaller  the 
mass  of  the  ion,  other  things  being  equal,  the  less  resist- 
ance it  will  offer  when  moving  through  the  solvent.  There- 
fore, rise  hi  temperature  should  increase  the  velocity  of 
the  ion. 

COEFFICIENTS  FOB  SLIGHTLY  HYDRATED  SALTS 
TABLE  A.  —  Temperature  coefficients  of  conductivity 


Substances  with  slight 
hydrating  power 

Temperature  coefficients  in 
conductivity  units 

25°  to  35° 

50°  to  65° 

0  =  8 

v  =  1024 

0  =  8 

v  =  1024 

Sodium  chloride  

2.00 
1.89 
2.12 
2.04 
2.39 
2.43 
2.38 
2.08 
2.04 
2.20 
2.42 
2.32 
2.17 

2.46 

2.54 
2.45 
2.84 
2.91 
2.91 
2.16 
2.31 
2.56 
2.94 
2.86 
2.50 

2.27 
2.18 
2.33 
2.02 
2.45 
2.45 
2.65 
2.31 
2.29 
2.34 
2.51 
2.58 
2.33 

2.82 
2.79 
3.14 
2.67 
3.11 
3.11 
3.37 
2.83 
2.23 

3^69 
3.11 
3.04 

Sodium  bromide 

Sodium  iodide 

Sodium  nitrate 

Potassium  chloride 

Potassium  bromide. 

Potassium  iodide.  . 

Potassium  nitrate.   . 

Potassium  permanganate  
Potassium  sulphocyanate  

Ammonium  chloride  

Ammonium  bromide  

Ammonium  nitrate  

If  decreasing  complexity  of  the  hydrate  formed  by  the  ion 
with  rise  hi  temperature  plays  any  prominent  part  in  deter- 
mining the  large  temperature  coefficients  of  conductivity, 
since  the  complexity  of  such  hydrates  would  decrease  more 
with  rise  hi  temperature,  we  should  expect  to  find  that  the 
ions  which  have  the  greatest  hydrating  power  would  have  the 
largest  temperature  coefficients  of  conductivity.  This  is  a  con- 
crete and,  it  would  seem,  necessary  consequence  of  the  theory 
of  hydrates  in  aqueous  solutions.  Further,  it  is  one  which 
can  be  tested  directly  by  experiment.  Is  it  true? 

We  have  seen  that  the  hydrating  power  of  a  salt,  or 
the  ions  into  which  it  dissociates,  is  approximately  pro- 
portional to  the  number  of  molecules  of  water  with  which 
it  crystallizes.  This  is  the  same  as  to  say  that  the  salt 


THE  NEWER  HYDRATE  THEORY 


319 


which  has  the  greatest  power  to  bring  water  with  it  out  of 
solution  is  the  one,  other  things  being  equal,  which  would 
hold  the  largest  number  of  molecules  of  water  hi  combina- 
tion with  it  hi  solution.  The  question  is,  therefore,  is 
there  any  relation  between  the  number  of  molecules  of 
water  with  which  a  salt  crystallizes  and  its  temperature 
coefficients  of  conductivity? 

This  relation  has  already  been  discussed  in  Publication 
No.  170  of  the  Carnegie  Institution  of  Washington,  hi 
which  the  results  of  the  work  on  conductivity  and  dis- 
sociation have  been  published.  Tables  A  and  B  showing 
temperature  coefficients  hi  conductivity  units  between  the 
temperatures  25°  and  35°,  on  the  one  hand,  and  between 
50°  and  65°  on  the  other,  at  the  dilutions  |  and  T^Vr 
normal,  are  taken  from  the  monograph  referred  to  above. 

COEFFICIENTS  FOB  STRONGLY  HYDRATED  SALTS 
TABLE  B.  —  Temperature  coefficients  of  conductivity 


Substances  with  large 
hydrating  power 

Temperature  coefficients  in 
conductivity  units 

25°  to  35° 

50°  to  65° 

v  =  S 

c  =  1024 

v  =  S 

0  =  1024 

Calcium  chloride 

3.49 
3.73 
3.09 
3.37 
3.66 
2.76 
3.09 
3.40 
3.55 
3.10 
3.13 
3.41 
3.21 
3.39 
3.32 
3.20 
3.18 
4.57 
4.19 

4.85 
5.00 
4.79 
5.13 
5.27 
4.86 
4.74 
4.72 
4.44 
4.78 
4.47 
5.04 
4.58 
4.95 
4.96 
4.67 
4.88 
8.64 
7.86 

4.03 
3.33 
3.92 
4.08 
3.58 
3.34 
3.61 
4.08 
3.57 
3.43 
3.61 

3.54 
3.75 
3.05 

5.16 

4.87 

6.03 
6.02 

5.41 

12.49 
11.65 

Calcium  bromide.  .    .      .        ... 

Calcium  nitrate  

Strontium  chloride  

Strontium  bromide  

Strontium  nitrate  

Barium  nitrate 

Magnesium  chloride 

Magnesium  bromide                   .    . 

Magnesium  nitrate  

Zinc  nitrate  

Nickel  chloride 

Nickel  nitrate 

Cobalt  chloride 

Cobalt  bromide 

Cobalt  nitrate 

Cupric  nitrate 

Aluminium  chloride  . 

Aluminium  nitrat.fi  T  -  -  ,  , 

320  THE  NATURE  OF  SOLUTION 

We  have  seen  that  the  hydrates  formed  by  a  large 
number  of  salts,  including  those  given  in  tables  A  and  B 
have  already  been  worked  out,1  and  that  water  of  crys- 
tallization is  a  rough  measure  of  water  of  hydration. 
The  salts  hi  table  A  crystallize  with  little  or  no  water, 
and  hi  aqueous  solution  are  very  little  hydrated;  those 
hi  table  B  hi  general  crystalhze  with  large  amounts  of 
water  and  are  strongly  hydrated  compounds. 

Let  us  compare  the  temperature  coefficients  of  conductiv- 
ity in  conductivity  units  (which  are  the  actual  increases 
in  molecular  conductivity  per  degree  rise  in  temperature) 
of  the  substances  in  table  A  with  those  hi  table  B.  It 
will  be  seen  that  the  coefficients  for  the  substances  in 
table  A  are,  at  all  dilutions  and  temperatures,  much  smaller 
than  those  hi  table  B.  In  making  this  comparison  we 
must,  of  course,  take  into  account  the  fact  that  the  sub- 
stances hi  table  A  are  binary  electrolytes,  each  molecule 
breaking  down  into  2  ions,  while  the  substances  recorded 
in  table  B  are  all  ternary  electrolytes,  each  molecule 
breaking  down  into  3  ions,  except  the  two  salts  of  alumin- 
ium which  are  quaternary  electrolytes,  each  molecule 
yielding  4  ions.  Even  taking  all  of  these  facts  into  account 
the  temperature  coefficients  of  conductivity  for  the  slightly 
hydrated  salts  are  much  smaller  than  those  for  the  strongly 
hydrated  compounds.  This  is  exactly  what  would  be 
expected.  The  complexity  of  the  hydrates  of  slightly 
hydrated  salts  could  change  only  a  little  with  rise  in  tem- 
perature. Consequently,  the  mass  of  the  hydrated  ion 
would  change  only  slightly  with  rise  in  temperature,  and 
this  effect  of  temperature  on  conductivity  would  be  very 
small. 

Relations  Between  the  Coefficients.  —  Another  relation 
manifests  itself  when  we  compare  the  results  in  table  A 
with  one  another,  and  those  in  table  B  with  one  another. 
If  the  temperature  coefficient  of  conductivity  is  a  function 
of  the  decreasing  complexity  of  the  hydrate  formed  by 

1  Carnegie  Institution  of  Washington,  Publication  No.  60  (1907). 


THE  NEWER  HYDRATE  THEORY         321 

the  ion,  as  the  temperature  is  raised  we  should  expect  that 
those  substances  which  have  equal  hydrating  power  would 
have  approximately  the  same  temperature  coefficients  of  con- 
ductivity. 

The  substances  in  table  A  have  only  slight  hydrating 
power,  shown  by  the  fact  that  they  crystallize  with  little 
or  no  water.  The  fact  is,  their  temperature  coefficients 
of  conductivity  are  all  of  the  same  order  of  magnitude. 

The  salts  hi  table  B  have  different  hydrating  powers, 
but  all  have  large  hydrating  power.  They  have  tempera- 
ture coefficients  of  conductivity  of  the  same  order  of  mag- 
nitude, with  the  exception  of  the  salts  of  aluminium. 

The  chloride  of  aluminium  crystallizes  with  6  molecules 
of  water  and  the  nitrate  with  8.  These  salts,  as  has 
already  been  pointed  out,  break  down  yielding  4  ions 
each.  Their  temperature  coefficients  are  larger  than  those 
of  the  ternary  electrolytes. 

The  more  dilute  the  solution,  the  more  complex  the 
hydrate  formed  by  the  molecule  or  the  ion.  This  is  but 
the  expression  of  the  action  of  mass;  the  more  water 
there  is  present,  the  more  will  be  combined  with  the  dis- 
solved substance.  The  more  complex  the  hydrate  the 
greater  the  change  in  the  complexity  of  the  hydrate  with 
rise  in  temperature.  Since  the  magnitude  of  the  tempera- 
ture coefficients  of  conductivity  seems  to  be  a  function  of 
the  change  hi  the  complexity  of  the  hydrate  with  rise  in 
temperature,  it  follows,  from  the  hydrate  theory,  that  the 
temperature  coefficients  of  conductivity  for  any  given  sub- 
stance should  be  greater  at  the  higher  dilution  than  at  the 
lower. 

A  comparison  of  the  results  at  the  two  dilutions  for 
any  given  substance  in  table  A  or  table  B  will  show  that 
the  above  consequence  of  the  hydrate  theory  is  confirmed 
by  the  facts.  The  temperature  coefficients  are  larger  at 
the  higher  dilution  for  every  substance  recorded  in  both 
tables. 

One  other  relation  should  be  pointed  out  before  leaving 


322  THE  NATURE  OF  SOLUTION 

the  discussion  of  the  temperature  coefficients  of  conduc- 
tivity. We  have  seen  that  the  hydrates  are  unstable, 
and  that  with  rise  in  temperature  they  break  down.  The 
higher  the  temperature  to  which  they  are  heated  the  more 
unstable  they  become.  We  should,  therefore,  expect  the 
hydrates  to  break  down  more  rapidly  as  the  temperature 
goes  higher.  If  this  were  the  case,  the  higher  the  tempera- 
ture of  the  solution,  the  larger  the  temperature  coefficients  of 
conductivity.  If  we  compare  the  results  for  any  given 
substance  in  table  A  or  B  we  will  find  that  such  is  the 
case.  The  temperature  coefficients  for  any  given  dilution 
are  higher  between  50°  and  65°  than  between  25°  and  35°. 

The  above  four  conclusions  from  the  solvate  theory  of 
solution,  as  far  as  aqueous  solutions  are  concerned,  are 
confirmed  at  every  point  by  the  results  of  measuring  the 
temperature  coefficients  of  conductivity.  Without  this 
theory  it  does  not  appear  to  be  simple  to  explain  the  above 
relations.  The  agreement  between  the  four  deductions 
from  the  theory  and  the  experimental  results  is  so  satis- 
factory that  it  is  regarded  as  strong  evidence  in  favor  of 
the  general  correctness  of  the  theory. 

Relation  Between  the  Hydration  of  the  Ions  and  Their 
Ionic  Volumes.  —  Jones  and  Pearce1  worked  out  the 
approximate  composition  of  the  hydrates  formed  by  a 
large  number  of  salts,  using  the  freezing-point  and 
conductivity  methods  already  referred  to.  They  found 
the  following  relation  between  the  volumes  of  the  ions 
and  then*  power  to  form  hydrates.  The  atomic  volume 
curve2  is  obtained  by  plotting  the  atomic  weights  of 
the  elements  as  abscissae  against  the  atomic  volumes  as 
ordinates.  This  curve,  as  is  well  known,  contains  well- 
defined  maxima  and  minima.  At  the  maxima  are  the  al- 
kali elements,  the  three  with  the  largest  atomic  volumes 
being  potassium,  rubidium  and  caesium.  The  salts  of 

1  Carnegie  Institution  of  Washington,  Publication  No.  180,  p.  57  (1913); 
Amer.  Chem.  Journ.,  38,  736  (1907). 

2  See  Elements  of  Physical  Chemistry,  by  the  author,  4th  edition,  p.  29. 
(The  Macmillan  Co.) 


THE  NEWER  HYDRATE  THEORY         338 

these  elements  generally  crystallize  without  water,  and 
therefore  have  very  little  hydrating  power  in  aqueous  solu- 
tion. The  approximate  hydration1  of  salts  of  potassium 
has  been  determined  by  the  method  usually  employed  and 
has  been  found  to  be  small. 

Some  of  the  salts  of  lithium  and  sodium  crystallize 
with  2  and  3  molecules  of  water,  and  these  have  been 
shown  to  have  some  hydrating  power.1  The  atomic 
volumes  of  lithium  and  sodium  are  much  smaller  than 
those  of  potassium,  rubidium,  and  caesium. 

Turning  from  the  maxima  of  the  curve  to  the  minima, 
at  the  mininum  of  the  third  section  of  the  curve  are  iron, 
cobalt,  nickel,  and  copper.  Salts  of  these  metals  crystallize 
with  large  amounts  of  water,  and  hi  aqueous  solution  they 
form  complex  hydrates. 

Aluminium  falls  at  the  second  minimum  of  the  atomic 
volume  curve,  having  a  somewhat  greater  atomic  volume 
than  iron.  The  salts  of  aluminium  crystallize  with  large 
amounts  of  water,  some  of  them  with  6  and  8  molecules. 
In  aqueous  solution  they  form  complex  hydrates.2  Barium 
has  the  largest  atomic  volume  of  members  of  its  group;  its 
salts  crystallize  without  water  or  some  with  2  molecules 
of  water.  Many  of  the  salts  of  calcium,  strontium,  and 
magnesium  crystallize  with  6  molecules  of  water.  Mag- 
nesium has  the  smallest  volume  of  any  element  of  this 
group;  it  has  been  found  to  have  the  greatest  hydrating 
power  of  any  member  of  the  group.  Strontium  has  a 
slightly  larger  atomic  volume  than  calcium  and  has  a 
somewhat  smaller  power  to  form  hydrates.  Taking  all  of 
the  facts  into  account,  it  would  seem  that,  other  things 
being  equal,  the  smaller  the  cation  the  greater  its  hydra- 
ting  power.  This  raises  the  question,  which  ion  is  it  that 
forms  the  Tiydrate?  Do  both  ions  form  hydrates?  If  so, 
which  has  the  greater  hydrating  power? 

The  different  salts  of  certain  metals  have  approxi- 
mately the  same  hydrating  power.  The  common  con- 

1  Carnegie  Institution  of  Washington,  Publication  No.  60  (1907) .         2  Ibid. 


324  THE  NATURE  OF  SOLUTION 

stituent  of  these  salts  is  of  course  the  cation,  the  anion 
varying  from  salt  to  salt.  This  would  indicate  that  it  is 
primarily  the  cation  which  conditions  the  hydrating  power 
of  a  salt.  Since  the  different  salts  of  the  same  metal  do 
not  all  have  the  same  hydrating  power,  it  seems  reasonable 
to  assume  that  the  anion  has  some  power  to  form  hydrates 
in  the  presence  of  water.  The  cation  is,  then,  the  chief 
hydrating  agent,  and  its  hydrating  power  seems  to  be  a 
function  of  its  size  or  atomic  volume  —  the  smaller  the 
ion  the  greater  its  power  to  hold  water  hi  combination 
with  it  hi  aqueous  solution. 

This  raises  the  question,  why  is  this  the  case?  It  has 
occurred  to  the  writer  that  the  electrical  density  of  the 
charge  on  the  ion  may  have  something  to  do  with  this 
relation.  Other  things  being  equal,  the  smaller  ion  has 
the  greater  density  of  charge  upon  its  surface;  this  might 
enable  it  to  hold  more  molecules  of  water  in  combination 
with  itself.  There  seem,  however,  to  be  certain  physical 
objections  to  this  explanation  of  the  relation  in  question. 
Whatever  the  explanation,  the  fact  remains. 

Hydration  of  the  Ions  and  the  Velocities  with  Which 
they  Move.  —  Certain  apparent  discrepancies  presented 
themselves  hi  the  velocities  of  the  different  ions,  which, 
for  a  time,  could  not  be  explained.  It  was  known  that  the 
lithium  ion,  under  the  same  driving  force,  moves  more 
slowly  than  potassium;  and  yet  it  has  smaller  volume  and 
smaller  mass.  It  was  not  until  it  was  shown1  that  the 
lithium  ion  is  more  strongly  hydrated  than  sodium  or 
potassium  that  this  fact  could  be  explained,  and  other 
apparent  discrepancies  presented  themselves.  A  relation 
between  the  migration  velocities  of  the  ions  and  their 
hydrating  power  was  worked  out  by  Jones  and  Pearce.2 
Their  discussion  is  repeated  here  to  bring  out  the  point  in 
question. 

The  velocities  of  the  ions  in  moving  through  any  given 

1  Carnegie  Institution  of  Washington,  Publication  No.  60  (1907). 

2  Ibid.,  Publication  No.  180,  pp.  84-86. 


THE  NEWER  HYDRATE  THEORY         325 

medium  may  be  assumed  to  vary  inversely  as  their  mass, 
the  driving  force  being  constant.  Their  velocities  would 
also  vary  inversely  as  their  volumes.  Mass  being  constant, 
we  should  expect  the  ions  with  the  smallest  atomic  volumes 
to  move  the  swiftest  under  a  constant  driving  force,  while 
the  facts  are  often  the  opposite.  Leaving  out  of  account 
the  hydrogen  and  hydroxyl  ions,  potassium,  rubidium,  and 
caesium  have  very  great  velocities  and  the  largest  volumes; 
while  the  ions  of  the  iron  and  copper  group  have  the  small- 
est volumes  and  very  small  velocities.  The  meaning  of 
this  apparent  discrepancy  can  be  seen  at  once  by  com- 
paring the  atomic  volume  curve  and  the  migration  velocity 
curve. 

The  elements  with  the  smallest  atomic  volumes  have  the 
greatest  hydrating  power.  When  these  atoms  appear  as 
ions  they  frequently  have  the  smallest  velocities.  There- 
fore, the  ions  with  smallest  velocities  have  the  greatest 
hydrating  power.  To  discuss  the  relations  somewhat  hi 
detail,  the  atomic  volumes  of  potassium,  rubidium,  and 
caesium  increase  rapidly  with  increasing  atomic  weight, 
and  their  salts  generally  crystallize  without  water.  The 
atomic  volumes  of  sodium  and  lithium  are  less  than  half 
that  of  potassium,  and  yet  their  ionic  velocities  are  only 
about  two-thirds  that  of  potassium.  It  will  be  recalled 
that  salts  of  sodium  and  lithium  may  crystallize  with  2 
or  3  molecules  of  water.  We  may  therefore  assume  that 
the  increase  hi  the  volume  and  mass  of  the  lithium  and 
sodium  ions,  due  to  the  formation  of  a  hydrate,  decreases 
the  velocity  of  these  ions  below  that  of  potassium. 

The  small  velocity  of  the  lithium  ion  was,  as  we  have 
seen,  for  a  long  tune  unexplained.  Lithium  has  a  volume 
only  about  half  that  of  sodium,  and  the  largest  ascertained 
amount  of  water  with  which  the  salts  of  lithium  crystallize 
is  3.  The  maximum  amount  for  many  of  the  salts  of 
sodium  is  2.  The  lithium  ion  is,  in  general,  more  hydrated 
than  the  sodium  ion,  and  its  velocity  is  therefore  decreased 
more  by  hydration.  Notwithstanding  its  smaller  volume 


326  THE  NATURE  OF  SOLUTION 

and  lighter  mass,  on  account  of  its  greater  hydration 
lithium  moves  with  about  the  same  velocity  as  sodium. 

The  calcium  atom  is  slightly  larger  than  sodium,  but  the 
calcium  ion  has  considerably  smaller  velocity.  This  is  un- 
doubtedly due  primarily  to  its  much  greater  hydrating 
power.  Within  this  group  the  atomic  volumes  increase  with 
increasing  atomic  weight.  The  velocities  of  calcium  and 
strontium,  with  many  salts  crystallizing  with  6  molecules  of 
water,  are  approximately  equal  to  that  of  barium.  Many  of 
the  salts  of  barium  crystallize  with  2  molecules  of  water 
or  water-free.  The  larger  mass  of  the  barium  atom  itself 
diminishes  the  velocity.  Magnesium,  with  about  half  the 
volume  of  calcium,  has  nearly  the  same  velocity,  due  to 
its  greater  hydrating  power.  The  cobalt,  nickel  and 
copper  atoms  have  nearly  the  same  volumes  and  approxi- 
mately the  same  hydrating  power.  Their  ions  have  ap- 
proximately the  same  velocities. 

The  atomic  volumes  of  chlorine,  bromine,  and  iodine 
are  approximately  the  same.  If,  as  ions,  they  hydrate 
at  all  we  should  expect  the  same  order  of  hydration  for 
all  three,  as  has  been  made  probable.  We  should  expect 
them  to  have  velocities  of  the  same  order  of  magnitude, 
and  such  is  the  fact. 

The  silver  ion  is  the  only  well-established  exception. 
Silver  has  a  small  volume  and  many  of  its  salts  crystallize 
without  water.  Although  it  has  small  volume,  it  appar- 
ently has  but  little  hydrating  power.  Notwithstanding  its 
considerable  mass,  with  its  small  volume  and  small  hydra- 
ting  power  we  should  expect  it  to  have  a  fairly  high  velocity. 
The  fact  is,  the  velocity  of  the  silver  ion  is  slightly  less 
than  that  of  chlorine,  bromine,  and  iodine. 

The  general  truth  of  the  relation  that  the  ions  with  the 
smallest  velocities  have  the  greatest  hydrating  power  is, 
then,  established  by  the  facts,  the  great  hydrating  power 
being  one  of  the  factors  conditioning  the  small  velocity. 


CHAPTER  XV 

THE   SOLVATE  THEORY  OF  SOLUTION 

Hydrate  Theory  for  Aqueous  Solutions  Becomes  the 
Solvate  Theory  for  Solutions  in  General.  —  The  earliest 
work  on  the  problem  of  the  nature  of  solution  was  limited 
to  water  as  the  solvent.  It  was  found  that  salts  in  general 
have  the  power  to  combine  with  more  or  less  of  the  water 
hi  which  they  are  dissolved  —  have  a  greater  or  less  hydra- 
ting  power.  This  power  is,  however,  possessed  to  a  very 
different  degree  by  the  different  compounds. 

It  having  been  made  probable  that  hydration  exists 
hi  aqueous  solution,  the  question  arose,  do  dissolved  sub- 
stances have  the  power  to  combine  with  other  solvents 
hi  which  they  are  dissolved? 

To  test  this  Jones  and  Getman1  studied,  by  the  boiling- 
point  method,  solutions  of  lithium  chloride  and  nitrate 
and  calcium  nitrate  in  ethyl  alcohol.  They  used  also  a 
number  of  other  salts.  It  was  found  that  the  molecular 
rise  in  the  boiling-point  was  not  only  greater  than  the 
theoretical  rise  at  nearly  all  of  the  concentrations  studied, 
but  the  molecular  rise  increases  rapidly  with  the  concen- 
tration of  the  solution.  The  molecular  rise  of  the  boil- 
ing-point of  ethyl  alcohol,  produced  by  lithium  chloride, 
increases  from  1.28°  at  0.07  normal  to  2.43°  at  2.07  normal. 
In  calculating  the  theoretical  molecular  rise  the  dissocia- 
tion is,  of  course,  taken  into  account.  The  dissociation 
decreases  with  the  concentration,  which  would  tend  to 
decrease  the  molecular  rise  in  the  boiling-point.  Notwith- 
standing this  influence,  we  have  seen  that  the  molecular 
rise  in  the  boiling-point  of  solutions  of  certain  salts  in 

1  Amer.  Chem.  Journ.,  32,  338  (1904). 


328  THE  NATURE  OF  SOLUTION 

ethyl  alcohol  increases  as  the  concentration  of  the  solution 
increases. 

The  differences  between  the  theoretical  and  the  experi- 
mental results  are  in  some  cases  quite  large.  Jones  and 
Getman1  interpreted  these  results  in  ethyl  alcohol  in  a 
manner  analogous  to  that  which  they  had  adopted  in  the 
case  of  aqueous  solutions.  The  abnormally  large  rise  in 
the  boiling-point  of  ethyl  alcohol,  produced  by  certain 
salts,  and  the  increase  in  the  molecular  rise  of  the  boiling- 
point  with  increase  in  the  concentration  of  the  solution, 
are  due  to  combination  between  the  dissolved  substance 
and  part  of  the  solvent  —  to  the  formation  of  alcoholates 
in  solution.  The  part  of  the  alcohol  that  is  combined 
with  the  dissolved  substance  is  thus  removed  from  the 
field  of  action  as  far  as  solvent  is  concerned.  There  being 
less  alcohol  present  acting  as  solvent,  the  rise  in  its  boiling- 
point  produced  by  a  given  amount  of  dissolved  substance 
would  be  larger  than  if  all  the  alcohol  were  playing  the 
part  of  solvent. 

Further,  if  the  dissolved  substance  combines  with  a 
part  of  the  alcohol,  the  more  concentrated  the  solution  the 
greater  the  total  amount  of  alcohol  held  in  combination. 
This  would  explain  the  increase  in  the  molecular  rise  in 
the  boiling-point  with  increase  in  the  concentration  of  the 
solution.  This  suggestion  of  combination  between  a  part 
of  the  solvent  and  the  dissolved  substance  explains  the 
facts  in  alcoholic  solutions  just  as  well  as  the  hydrate 
theory  explains  the  facts  in  aqueous  solutions. 

The  work  of  Jones  and  Getman  with  solutions  in  ethyl 
alcohol  as  the  solvent  was  extended  by  Jones  and  McMas- 
ter  to  methyl  alcohol.  They  also  extended  the  work  in 
ethyl  alcohol  as  the  solvent  and  repeated  a  part  of  the 
work  of  Jones  and  Getman  obtaining  results  of  the  same 
general  character  as  had  been  found  by  the  earlier  workers. 

They  used  the  boiling-point  method  with  methyl  alcohol 
as  the  solvent,  and  the  chloride,  bromide  and  nitrate  of 

1  Amer.  Chem.  Jour.,  32,  339  (1904). 


SOLVATE  THEORY  OF  SOLUTION  329 

lithium  as  the  dissolved  salts.  The  molecular  rise  in  the 
boiling-point,  even  hi  the  most  dilute  solutions,  was  greater 
than  could  be  accounted  for  by  the  dissociation.  This  is, 
of  course,  entirely  incapable  of  accounting  for  the  increase 
in  the  molecular  rise  with  increase  in  the  concentration  of 
the  solution,  which  manifests  itself  hi  the  case  of  every 
salt  studied  hi  this  solvent,  dissociation  decreasing  with 
increase  hi  concentration,  which  would  tend  to  diminish 
the  molecular  rise  in  the  boiling-point. 

The  magnitude  of  the  molecular  rise  hi  the  most  con- 
centrated solutions  is  very  large  indeed.  It  is  almost 
twice  the  boiling-point  constant,  or  normal  molecular  rise 
for  this  solvent;  and  the  dissociation  of  such  solutions  is 
certainly  not  more  than  25  to  30  per  cent,  and  probably 
less  than  this  value. 

These  results  were  interpreted  as  were  those  hi  ethyl 
alcohol  as  the  solvent  —  there  is  combination  between  a 
part  of  the  alcohol  present  and  the  dissolved  substance, 
forming  methyl  alcoholates.  As  the  concentration  in- 
creases, more  and  more  alcohol  is  held  hi  combination  by 
the  dissolved  substance;  consequently,  there  is  an  increase 
in  the  molecular  rise  of  the  boiling-point. 

It  thus  seems  that  evidence  was  furnished  of  combina- 
tion between  methyl  alcohol  and  the  dissolved  substance, 
on  the  one  hand,  and  ethyl  alcohol  and  the  dissolved 
substance  on  the  other.  As  we  shall  see  later,  evidence 
has  been  obtained  of  combination  between  acetone  and 
substances  dissolved  hi  it;  and  other  solvents  have  been 
brought  within  the  scope  of  this  work. 

In  every  case  thus  far  investigated  there  seems  to  be 
good  evidence  in  favor  of  the  view  that  there  is  combina- 
tion between  the  dissolved  substance  and  a  part  of  the 
solvent  present.  In  a  word,  combination  of  solvent  with 
dissolved  substance  —  solvation  —  seems  to  be  a  more  or 
less  general  phenomenon.  The  original  hydrate  theory  thus 
becomes  the  solvate  theory  of  solution. 


330  THE  NATURE  OF  SOLUTION 

SPECTROSCOPIC  EVIDENCE  BEARING  ON  THE  SOLVATE 
THEORY  OF  SOLUTION 

Work  of  Uhler.  —  Work  on  the  absorption  spectra  of 
solutions  has  been  in  progress  in  the  author's  laboratory 
continuously  for  nine  years.  This  work  was  undertaken 
in  connection  with  its  bearing  on  the  solvate  theory  of 
solution.  What  connection  is  there  between  solvation 
and  the  power  of  solutions  to  absorb  light? 

It  is  well  known  that  absorption  of  light  means  that  the 
wave-lengths  of  light  set  something  vibrating  with  periods 
the  same  as  their  own.  Selective  absorption  of  light  or  the 
absorption  of  certain  wave-lengths  of  light  means  that 
the  wave-lengths  absorbed  set  something  vibrating  with 
their  own  periods.  Absorption  of  light  is,  then,  a  reso- 
nance phenomenon.  Absorption  of  light  by  a  dissolved 
substance  means  that  something  in  the  solution  must  be 
thrown  into  resonance  with  the  light  —  must  be  set  vibra- 
ting with  the  same  periods  as  the  light-waves.  Many 
dissolved  substances  absorb  certain  wave-lengths  only. 
This  means  that  those  particular  wave-lengths  of  light  find 
something  in  the  solution  which  they  can  set  vibrating 
with  their  own  periods.  Transparency  means  lack  of 
resonance,  opacity  means  resonance.  The  color  of  any 
given  solution  is  determined  by  the  wave-lengths  of  light 
which  are  not  absorbed.  A  red  solution  is  one  which 
allows  the  long  wave-lengths  to  pass  through.  A  blue 
solution  is  one  which  allows  the  short  wave-lengths  to 
pass  through.  That  particle  in  solution  which  is  thrown 
into  resonance  by  the  light  is  called  the  resonator.  This 
was  formerly  supposed  to  be  the  molecule  or  the  ion,  but 
is  now  thought  to  be  the  electron.  Whatever  the  nature 
of  the  resonator,  the  absorption  of  light  by  dissolved  sub- 
stance is  due  to  it. 

The  line  of  thought  which  led  us  to  take  up  the  study 
of  the  absorption  spectra  of  solutions  in  connection  with 
the  solvate  theory  of  solution  is  as  follows:  The  absorp- 


SOLVATE  THEORY  OF  SOLUTIONS  331 

tion  of  light  being  due  to  a  resonator,  this  would  have 
different  resonance  when  anhydrous  than  when  combined 
with  molecules  of  the  solvent.  In  general,  the  resonance 
would  be  different  when  the  resonator  was  unsolvated 
than  when  it  was  solvated.  The  color  of  the  solution 
being  due  to  the  resonator,  the  solution  could  reasonably 
be  expected  to  have  different  color  when  the  resonator 
was  solvated  than  when  it  was  unsolvated.  The  study  of 
the  color  of  solutions,  and  the  changes  hi  the  color  when  the 
resonator  underwent  changes  hi  solvation,  might  give  some 
clue  to  the  changes  in  solvation. 

It  is  a  comparatively  simple  matter  to  change  solvation 
hi  solution;  it  is  only  necessary  to  change  the  concentra- 
tion of  the  solution.  The  more  dilute  the  solution  the 
more  complex  the  solvates  formed.  We  shall  see  that  this 
often  produces  a  marked  change  hi  the  absorption  spectra. 
We  can  diminish  the  complexity  of  the  solvates  by  raising 
the  temperature.  Frequently  this  also  produces  marked 
changes  in  the  absorption.  Addition  of  a  dehydrating 
agent  will  change  the  hydration  of  any  given  salt.  This 
frequently  changes  the  absorption  spectra  and  the  color 
of  a  solution;  and  there  are  many  other  ways  of  changing 
solvation.  These  frequently  produce  concomitant  changes 
in  the  absorption  spectra. 

A  salt  dissolved  hi  water  may  form  hydrates,  hi  alcohol 
alcoholates,  in  acetone  acetonates,  in  glycerol  glycerolates, 
etc.  We  should  expect  these  different  solvates  to  affect 
the  resonator  or  resonators  differently.  We  shall  see  that 
this  is  true. 

With  this  idea  in  mind,  work  was  begun  in  this  labora- 
tory on  the  study  of  the  absorption  spectra  of  solutions. 
The  first  investigation  was  carried  out  by  Dr.  Uhler  and 
the  writer.  The  work  consisted  largely  in  devising  a 
method  and  apparatus  for  studying  the  property  of  solu- 
tions to  absorb  light.  The  key  to  the  method  consisted  hi 
using  a  grating  instead  of  a  prism  spectroscope.  This  gave 
much  greater  dispersion,  and  brought  out  many  new  lines 


332  THE  NATURE  OF  SOLUTION 

and  bands.  A  form  of  cell  was  devised  for  holding  solu- 
tions in  non-aqueous  solvents  which  avoided  the  use  of 
all  cement.  The  details  of  this  phase  of  the  work  are  all 
given  hi  Publication  No.  60  of  the  Carnegie  Institution  of 
Washington.  The  effect  on  the  absorption  spectra  of  in- 
creasing the  concentration  of  the  solution  was  studied  and 
it  was  found  that,  hi  general,  the  effect  of  increasing  the 
concentration  of  the  solution  was  to  widen  the  absorption 
bands.  As  the  solvates  became  simpler  the  absorption 
bands  became  broader. 

Another  method  of  simplifying  the  hydrates  existing 
in  an  aqueous  solution  was  to  add  a  dehydrating  agent 
hi  the  form  of  a  second  salt.  It  was  found  that  this  also 
produced  a  widening  of  the  absorption  bands.  This  was 
in  keeping  with  the  effect  of  increasing  the  concentration 
of  the  solution,  which  also  simplified  the  hydrates. 

Jones  and  Uhler  also  studied  the  effect  of  adding  water 
to  solutions  in  non-aqueous  solvents.  Thus,  water  was 
added  to  solutions  in  methyl  and  ethyl  alcohols,  acetone, 
etc.  The  effect  of  adding  water  was  to  narrow  the  ab- 
sorption bands.  All  of  these  results  were  regarded  as  in 
keeping  with  the  solvate  theory  of  solution. 

"Solvent  Bands":  Work  of  Anderson.  —  The  work  of 
Jones  and  Uhler  on  the  absorption  spectra  of  solutions 
was  greatly  extended  in  a  number  of  directions  by  Jones 
and  Anderson.1  They  worked  with  salts  of  cobalt, 
nickel,  copper,  iron,  chromium,  neodymium,  praseodymium 
and  erbium.  Only  that  phase  of  the  work  will  be  dis- 
cussed here  which  bears  most  directly  on  the  solvate 
theory  of  solution.  The  feature  of  this  work,  which  bears 
most  directly  on  the  solvate  theory  of  solution,  came  out 
as  the  result  of  studying  the  absorption  spectra  of  solu- 
tions of  salts  of  neodymium  and  praseodymium,  especially 
of  neodymium. 

Neodymium  chloride  was  found  to  have  quite  different 

1  Carnegie  Institution  of  Washington,  Publication  No.  110  (1909).  Amer. 
Chem.  Journ.,  41,  163  (1909). 


SOLVATE  THEORY  OF  SOLUTION  333 

absorption  in  water  from  what  it  had  in  methyl  alcohol. 
This  made  it  desirable  to  study  the  absorption  spectrum 
of  this  salt  hi  mixtures  of  methyl  alcohol  and  water.  By 
changing  the  composition  of  the  mixtures  of  the  two  sol- 
vents, one  could  see  how  the  spectra  corresponding  to  the 
two  solvents  would  change. 

It  was  found  that  when  the  proper  mixture  of  alcohol 
and  water  was  used,  the  two  spectra  (the  one  corresponding 
to  the  alcoholic  solution  and  the  other  to  the  aqueous  solu- 
tion) coexisted  on  the  plate.  When  the  amount  of  water 
hi  the  mixed  solvents  increased,  the  "water  spectrum" 
came  out  more  strongly;  when  the  amount  of  alcohol 
present  was  increased,  the  "alcohol  spectrum"  came  out 
more  strongly.  When  the  amount  of  water  present 
exceeds  15  or  20  percent,  we  have  only  the  "water  spec- 
trum." As  the  amount  of  water  is  still  further  decreased 
by  the  addition  of  more  alcohol,  the  spectrum  consists 
of  the  "water  spectrum"  and  the  "alcohol  spectrum" 
superposed.  As  the  amount  of  water  is  diminished  below 
15  percent,  the  intensity  of  the  "water  spectrum"  be- 
comes less  and  less  and  the  intensity  of  the  "alcohol 
spectrum"  greater  and  greater. 

A  question  of  importance  hi  the  present  connection  is 
this:  Does  the  "water  spectrum"  gradually  change  over 
into  the  "alcohol  spectrum"  as  the  amount  of  alcohol 
present  is  increased,  or  do  we  have  here  two  separate  and 
distinct  spectra,  the  one  corresponding  to  the  aqueous 
solution,  and  the  other  to  the  alcoholic? 

To  test  this  point,  fairly  dilute  solutions  of  neodymium 
chloride  in  water,  in  methyl  alcohol,  and  hi  mixtures 
of  water  and  methyl  alcohol  were  used.  The  object  hi 
using  dilute  solutions  was  to  be  able  to  study  the  struc- 
ture of  the  bands  in  the  different  solvents.  In  the  more 
dilute  solutions  the  several  parts  of  any  given  band  would 
come  out  clearly  and  could  be  measured.  The  result  was 
to  show  that  the  "alcohol  spectrum"  was  quite  different 
from  the  "water  spectrum."  It  had  different  components 


334  THE  NATURE  OF  SOLUTION 

and  they  were  arranged  in  a  different  way  within  the 
bands. 

In  mixed  solvents,  then,  the  two  spectra  coexisted,  and 
the  one  did  not  pass  over  into  the  other  as  the  com- 
position of  the  mixture  of  alcohol  and  water  was  changed. 
The  "water"  spectrum  and  "methyl-alcohol"  spectrum 
had  equal  intensities  when  the  mixture  of  the  water  and 
methyl  alcohol  contained  from  6  to  8  percent  of  water. 

Neodymium  nitrate  shows  change  in  the  spectra  analo- 
gous to  those  manifested  by  the  chloride,  when  dissolved 
in  mixtures  of  water  and  one  of  the  non-aqueous  solvents. 
The  change  with  the  nitrate  is  not  so  striking  as  with  the 
chloride. 

Praseodymium  chloride  in  mixtures  of  water  and  methyl 
alcohol  shows  the  same  general  features  as  were  mani- 
fested by  the  chloride  of  neodymium.  In  the  case  of 
praseodymium  chloride  there  is  this  additional  feature:  in 
the  alcoholic  solution  an  entirely  new  band  appears,  having 
no  analogue  in  the  aqueous  solutions.  This  new  band 
in  the  ultra-violet  is  by  far  the  most  intense  in  the  entire 
spectrum  of  praseodymium  chloride.  On  adding  water 
to  the  alcoholic  solution  this  band  entirely  disappears. 
In  this  case  the  alcohol  spectrum  is  quite  different  from 
the  water  spectrum. 

These  results  show  beyond  question  that  the  solvent 
plays  an  important  role  in  the  absorption  of  light  by  solu- 
tions. The  question  arises,  what  is  this  r61e?  It  is 
difficult,  not  to  say  impossible,  to  explain  the  action  of  the 
solvent  on  any  other  ground  than  that  a  part  of  the  solvent 
combines  with  the  ions  and  molecules  of  the  dissolved 
substance,  the  solvated  parts  having  different  resonance 
from  the  unsolvated.  This  means  that  they  would  absorb 
different  wave-lengths  of  light.  The  alcoholates  would 
have  different  resonance  from  the  hydrates,  whence  the 
different  absorption  spectrum  in  alcohol  from  that  in  water. 

This  is  important  evidence  in  favor  of  solvation  in  solu- 
tion, and,'  as  we  shall  see,  many  examples  of  "solvent" 


SOLVATE  THEORY  OF  SOLUTION  335 

bands  were  brought  to  light  in  the  investigation  which 
followed. 

"Solvent  Bands":  Work  of  Jones  and  Strong.  — The 
work  of  Jones  and  Anderson  was  continued  by  Jones  and 
Strong.1  They  investigated  a  number  of  problems,  in- 
cluding the  effect  of  the  solvent  on  the  absorption  of  light 
by  the  dissolved  substance.  Jones  and  Anderson,  as  we 
have  just  seen,  had  found  one  good  example  of  the  exis- 
tence of  "  solvent  bands "  hi  the  absorption  spectra  of 
neodymium  and  praseodymium  salts  in  water  and  the 
alcohols.  The  question  arose,  was  this  a  phenomenon 
peculiar  to  these  salts,  or  does  the  solvent  play  a  general 
role  in  the  absorption  of  light  by  solutions? 

Jones  and  Strong  attempted  to  answer  this  question  by 
studying  a  large  number  of  salts  in  a  large  number  of 
solvents.  They  worked  especially  with  salts  of  neodym- 
ium and  uranium,  because  these  substances  had  sharp 
absorption  lines  and  bands  whose  positions  could  easily 
be  determined  with  reasonable  accuracy.  Work  was  done 
not  only  with  uranyl  salts,  but  with  uranous.  A  con- 
venient method  was  found  for  reducing  uranyl  salts  to 
the  uranous  condition,  and  uranous  salts  were  found  to 
have  very  sharp  absorption  lines. 

Absorption  Spectra  of  Uranium  Compounds.  —  Uranyl 
chloride  was  studied  in  the  following  solvents:  water; 
methyl,  ethyl,  propyl,  isopropyl,  butyl,  and  isobutyl  alco- 
hols; glycerol,  ether,  methyl  acetate,  and  formamide.  A 
comparison  of  the  wave-lengths  of  the  absorption  lines  and 
bands  hi  these  different  solvents  brought  out  the  fact 
that  the  wave-lengths  of  some  of  the  lines  and  bands 
differed  considerably  hi  the  different  solvents.  The  results 
here  showed  that  the  solvent  unquestionably  has  much 
to  do  with  the  absorbing  power  of  the  solution,  "solvent 
bands"  appearing  very  frequently.  The  wave-lengths 

1  Carnegie  Institution  of  Washington,  Publications  Nos.  130  (1910) ;  160 
(1911).  Amer.  Chem.  Journ.,  43,  37,  224  (1910);  46,  1(1910);  47,  27  (1912). 
Phys.  Zeit.  10,  499  (1909).  Phil.  Mag.,  19,  566  (1910).  Journ.  Chim.  Phys.t 
8,  131  (1910). 


336  THE  NATURE  OF  SOLUTION 

of  a  few  of  the  different  lines  and  bands  of  uranyl  chloride 
in  the  above-named  solvents  have  been  tabulated,1  and 
the  table  is  here  reproduced.  It  shows  at  a  glance  the 
different  wave-lengths  of  the  several  lines  and  bands  com- 
pared. 

Wave-lengths  of  uranyl  chloride  absorption  lines 

In  water XX  4025,  4170,  4315,  4460,  4560,  4740,  and  4920 

In  methyl  alcohol...  XX  4090,  4220,  4345,  4465,  4590,  4760,  and  4930 

In  ethyl  alcohol XX  4100,  4250,  ....   4400,  4580,  4750,  and  4900 

In  propyl  alcohol...  XX  4100,  4230,  ....   4400,  4580,  4750,  and  4910 

In  isopropyl  alcohol.  XX  4100,  4250,  ....   4360,  4560,  4750 

In  butyl  alcohol..  ..  XX  4100,  4240,  ....   4390,  4560,  4750,  and  4970 

In  isobutyl  alcohol XX4400,  4560,  4720,  and  4900 

In  ether XX  4040,  4160,  4300,  4444,  and  4630 

In  methyl  acetate. .  XX  4030,  4160,  4280,  4440,  4620,  4790,  and  4920 

Inglycerol XX  4025,  4140,  4260,  4400,  4540,  4720,  and  5050 

Informamide XX  4450,  4650,  and  4840 

The  absorption  spectra  of  uranyl  nitrate  in  mixtures  of 
water  and  methyl  alcohol  were  studied.  The  absorption 
in  water  was  much  less  than  in  pure  methyl  alcohol.  The 
addition  of  water  to  the  alcoholic  solution  diminished  the 
absorption.  In  the  mixtures  of  water  and  methyl  alcohol 
the  absorption  bands  became  very  broad.  A  study  of 
these  broadened  bands  shows  that  they  were  the  "  alcohol  " 
and  "  water  "  bands  coexisting,  and  that  one  set  of  bands 
was  not  simply  the  other  set  shifted  in  position.  The 
importance  of  this  fact  has  already  been  referred  to  in  the 
work  of  Jones  and  Anderson.  It  shows  that  the  "alcohol" 
bands  are  fundamentally  different  from  the  "water" 
bands.  Further,  the  intensity  of  the  solvent  bands  is  a 
function  of  the  relative  amounts  of  the  solvents  that  are 
present  in  the  mixture.  This,  as  has  been  pointed  out, 
indicates  the  existence  of  hydrates  in  the  aqueous  solutions 
and  of  alcoholates  in  solutions  in  alcohol,  these  solvates 
having  definite  resonance  and,  therefore,  definite  absorption 
spectra. 

One  of  the  most  striking  examples  of  solvent  bands  is 
shown  by  the  absorption  spectra  of  uranous  chloride  and 
bromide  in  a  mixture  of  water  and  methyl  alcohol.  We 

1  Journ.  Franklin  Inst.,  176,  528  (1913);  also  Phil  Mag.,  23,  730  (1912). 


SOLVATE  THEORY  OF  SOLUTION  337 

find  two  entirely  distinct  spectra,  one  belonging  to  each 
solvent.  Some  lines  and  bands  appear  in  the  one  solvent 
which  are  entirely  absent  from  the  other,  and  practically 
all  the  lines  and  bands  have  very  different  positions  hi  the 
two  solvents.  To  see  how  differently  the  spectra  appear, 
reference  must  be  made  to  plate  23  of  Publication  No.  160 
of  the  Carnegie  Institution  of  Washington. 

The  spectrum  of  uranous  chloride  in  water  is  not  only 
different  from  the  spectrum  hi  methyl  alcohol,  but  these 
are  both  different  from  the  spectrum  hi  acetone.  If  we 
compare  the  spectra  of  this  salt  hi  the  three  solvents,  we 
might  easily  conclude  that  we  were  dealing  with  three 
fundamentally  different  spectra,  and  the  only  change  is  in 
the  nature  of  the  solvent. 

Uranous  salts  hi  solvents  other  than  the  above  also 
show  very  characteristic  " solvent"  bands.  When  ethyl 
alcohol  is  added  to  an  aqueous  solution  of  uranous  chloride, 
a  marked  change  is  produced  hi  the  spectrum.  The 
"ethyl  alcohol"  bands  are  quite  different  from  the  "water" 
bands.  The  alcohol  bands,  or  the  water  bands,  can  be 
made  the  more  intense  by  simply  varying  the  relative 
proportions  of  the  two  solvents.  The  addition  of  acetone 
to  an  aqueous  or  methyl  alcohol  solution  of  uranous 
chloride  produces  a  marked  change  hi  the  spectrum.  A 
number  of  acetone  bands  appear,  these  being  different  from 
the  "  water  "  bands  on  the  one  hand,  and  from  the  "  alco- 
hol" bands  on  the  other. 

Uranous  chloride  dissolved  in  methyl  alcohol  has  an 
absorption  spectrum  very  similar  to  that  hi  ethyl  alcohol. 
This  would  be  expected,  on  account  of  the  close  similarity 
of  methyl  alcohol  and  ethyl  alcohol.  The  methyl  alcohol 
bands  are  of  slightly  shorter  wave-lengths. 

The  absorption  spectra  of  uranous  chloride  in  glycerol 
and  hi  mixtures  of  glycerol  and  water  were  also  studied. 
A  number  of  "glycerol"  bands  manifested  themselves, 
the  glycerol  absorption  being  very  different  from  that  of 
water. 


338  THE  NATURE  OF  SOLUTION 

The  absorption  spectrum  of  uranous  chloride  in  methyl 
alcohol  and  ether  was  also  studied.  The  solution  in 
methyl  alcohol  showed  complete  absorption  in  the  ultra- 
violet to  wave-length  X  3700,  while  the  addition  of  ether 
extended  the  absorption  to  X  3800.  The  addition  of  the 
ether  caused  the  absorption  to  shift  towards  the  red,  the 
magnitude  of  this  shift  being  from  10  to  30  A.U. 

It  has  already  been  pointed  out  that  salts  of  neodym- 
ium  are  especially  well  adapted  to  the  study  of  "solvent" 
bands,  on  account  of  the  sharpness  of  the  neodymium 
lines  and  bands,  and  the  accuracy  with  which  they  can  be 
measured.  Neodymium  salts  were  studied  in  a  number 
of  solvents,  and  a  few  of  the  results  obtained  are  given 
below.1 

Absorption  Spectra  of  Neodymium  Chloride  in  Various 
Solvents.  —  The  following  results  for  the  «  group  of 
absorption  bands,  lying  in  the  region  X  3400  to  X  3600 
obtained  with  noedymium  chloride  show  the  effect  of 
the  solvent  on  the  absorption  spectra  of  solutions  of  this 
compound.  The  bands  of  the  different  solvents  have  dif- 
ferent wave-lengths  and  different  relative  intensities. 

Having  found  that  the  solvent  played  an  important 
part  in  determining  the  absorption  of  light  by  the  dissolved 
substances,  Jones  and  Strong  used  isomeric  organic  sol- 
vents, to  see  whether  such  closely  related  compounds  would 
affect  differently  the  power  of  substances  dissolved  in 
them  to  absorb  light.  They  prepared  solutions  of  neodym- 
ium chloride  in  propyl  and  isopropyl  alcohols,  and  in 
butyl  and  isobutyl  alcohols,  and  photographed  the  absorp- 
tion spectra  of  this  salt  in  these  isomeric  solvents.  The 
results  show  different  absorption  lines  and  bands  in  the 
isomeric  solvents. 

If  we  compare  carefully  the  spectra  of  neodymium 
chloride  in  butyl  and  isobutyl  alcohols,  we  find  that  the 
bands  are  weak  and  diffuse  in  isobutyl  alcohol,  and  have 

1  See  Phil.  Mag.,  23,  737  (1912),  from  which  the  few  following  pages 
are  taken;  also,  Journ.  Franklin  InsL,  176, 531  (1913). 


SOLVATE  THEORY:  OF  SOLUTION 


339 


different  relative  intensities  from  what  they  have  in  the 
butyl  alcohol.    The  bands  in  butyl  alcohol  are  very  much 

Absorption  spectra  of  neodymium  chloride  in  certain  solvents 
a  GROUP 


In  water 

In  methyl 
and  ethyl 
alcohols 

In  propyl 
alcohol 

In 

isopropyl 
alcohol 

In  butyl 
alcohol 

In  isobuty] 
alcohol 

In 
glycerol 

XX 

XX 

XX 

XX 

XX 

XX 

XX 

3390 

3475 

3545 

3460 

3450 

3455 

3520 

3465 

3505 

3460 

3510 

3460 

3485 

3475 

3505 

3490 

3535 

3492 

3515 

3550 

3540 

3560 

3510 

3535 

3545 

3560 

.... 

3525 

.... 

3545 

3570 

.... 

.... 

3540 

.... 

3560 

.... 

.... 

3560 

.... 

3580 









finer  and  sharper  than  they  are  hi  isobutyl  alcohol.  Further, 
the  bands  of  neodymium  chloride  hi  isobutyl  alcohol  have 
slightly  greater  wave-lengths  than  hi  butyl  alcohol. 

Similar  results  were  obtained  with  /?,  y,  5  and  c  groups 
of  bands,  but  for  these  reference  must  be  had  to  Publication 
of  the  Carnegie  Institution  of  Washington,  No.  210  (1915). 

Absorption  Spectra  of  Neodymium  Nitrate  in  Different 
Solvents.  —  To  eliminate  the  possibility  of  the  effect  of 
the  solvent  on  absorption  spectra  being  due  to  anything 
inherent  in  the  nature  of  neodymium  chloride,  the  nitrate 
of  neodymium  was  studied  in  the  same  way  as  the  chloride. 

The  absorption  spectra  of  neodymium  nitrate  hi  water, 
in  methyl  alcohol,  hi  ethyl  alcohol,  hi  mixtures  of  these 
alcohols  and  water,  hi  propyl  and  isopropyl  alcohols,  hi 
butyl  and  isobutyl  alcohols,  hi  acetone  and  hi  mixtures  of 
acetone  and  water,  in  ethyl  acetate  and  in  formamide,  were 
carefully  photographed  and  studied.  Results  are  given  below 
in  the  case  of  neodymium  nitrate  only  for  the  a  bands. 

In  water.  —  Practically  the  same  as  the  bands  of  neodymium  chloride,  but 
the  bands  of  the  nitrate  are  broader  and  hazier  than  those  of  the  chloride. 

In  methyl  and  ethyl  alcohols.  —  There  are  only  two  bands  in  the  a  group, 
X  3465  and  X  3545. 


340  THE  NATURE  OF  SOLUTION 

In  propyl  alcohol.  —  XX  3455,  3500,  and  3585. 

In  isopropyl  alcohol.  —  XX  3460,  3505,  and  3535. 

In  butyl  alcohol  —  XX  3450,  3500,  and  3540. 

In  isobutyl  alcohol.  —  Ultraviolet  absorption  was  so  great  that  on  the 
plate  taken  the  a  group  did  not  appear.  The  absorption  in  general  is  the 
same  as  that  of  the  chloride  in  this  alcohol. 

In  acetone.  —  XX  3475,  and  3555. 

In  ethyl  acetate.  —  XX  3455,  3500,  and  3540. 

The  other  groups  of  absorption  bands  of  neodymium 
nitrate  hi  the  different  solvents  show  differences  hi  the 
wave-lengths  comparable  with  the  above;  but  these  re- 
sults suffice  to  show  the  effect  of  the  solvent  on  the  power 
of  neodymium  nitrate  to  absorb  light. 

The  above  is  strong  evidence  that  the  solvent  plays  an 
important  part  hi  the  absorption  of  light  by  substances 
dissolved  in  it.  When  we  take  into  account  the  number 
of  salts  studied  and  the  number  of  solvents  employed,  the 
evidence  is  little  short  of  proof.  The  only  reasonable  ques- 
tion is,  how  are  we  to  interpret  these  facts?  Before 
attempting  to  answer  this  question  we  should  take  into 
account  also  the  f ollowing  fact :  A  salt  dissolved  hi  a  given 
solvent  is  characterized  by  a  definite  absorption  spectrum. 
When  a  salt  is  dissolved  hi  mixtures  of  varying  propor- 
tions of  two  solvents,  only  two  definite  absorption  spectra 
appear,  one  being  characteristic  of  each  solvent.  One 
spectrum  does  not  gradually  change  into  the  other  as  the 
composition  of  the  mixed  solvent  changes,  but  only  the 
relative  intensities  of  the  two  spectra  vary.  Starting  with 
that  mixture  of  the  two  solvents  hi  which  both  of  the  spec- 
tra are  equally  intense,  if  we  diminish  the  amount  of  a 
relative  to  6,  the  spectrum  corresponding  to  a  becomes 
feebler  and  feebler,  and  the  spectrum  corresponding  to  b 
more  and  more  intense.  This  fact  was  first  noted  by 
Jones  and  Anderson,  and  was  repeatedly  confirmed  by 
the  work  of  Jones  and  Strong.  They  found  that  when 
neodymium  chloride  was  dissolved  in  a  mixture  of  methyl 
alcohol  and  water,  it  showed  a  definite  set  of  "water" 
bands  and  a  definite  set  of  "methyl  alcohol"  bands.  As 


SOLVATE  THEORY  OF  SOLUTION  341 

the  amount  of  water  in  the  solution  was  decreased  relative 
to  the  alcohol,  the  "water"  bands  decreased  in  intensity 
but  remained  in  the  same  position.  As  the  amount  of 
alcohol  in  the  solution  was  decreased  relative  to  the  water, 
the  "alcohol"  bands  decreased  hi  intensity,  but  their 
position  remained  unchanged. 

Jones  and  Anderson  interpreted  these  facts  as  strong 
evidence  in  favor  of  the  view  that  there  are  definite  hy- 
drates and  definite  alcoholates  hi  the  solution. 

The  spectroscopic  evidence  for  solvation  hi  solution 
furnished  by  Jones  and  Anderson  has,  as  has  already  been 
stated,  been  increased  many  fold  by  the  work  of  Jones 
and  Strong.  A  large  number  of  solvents  and  a  fairly 
large  number  of  salts  have  been  used,  and  the  existence 
of  solvent  bands  hi  general  has  been  established. 

The  question  of  the  relative  stability  of  the  different 
solvates  with  respect  to  various  physical  and  chemical 
agents  has  been  studied  at  length  by  Jones  and  Strong  by 
means  of  absorption  lines  and  bands.  It  would  lead  us 
beyond  the  scope  of  this  book  to  discuss  these  results 
hi  detail.  Suffice  it  to  say  that  the  hydrates  hi  general 
•are  the  most  persistent  of  all  the  solvates,  although  this 
depends  upon  the  conditions  to  which  the  solution  is  sub- 
jected. 

Transparency  of  Free  and  of  Combined  Water:  Work 
of  Guy.  —  The  work  on  the  absorption  spectra  of  solu- 
tions, at  the  time  that  Jones  and  Guy  began  their  inves- 
tigation, had  been  extended  to  between  6,000  and  7,000 
solutions.  In  all  of  this  work  the  grating  spectroscope 
had  been  used,  and  the  results  recorded  on  a  photo- 
graphic plate.  The  photographic  method  recorded  the 
positions  of  the  various  absorption  lines  and  bands,  but 
gave  only  a  qualitative,  or  at  best  a  roughly  quantitative, 
indication  of  the  relative  intensities  of  the  various  lines 
and  bands.  The  photographic  method  is,  generally  speak- 
ing, a  qualitative  method. 

If  we  are  ever  to  discover  relations  of  fundamental 


342  THE  NATURE  OF  SOLUTION 

significance  between  the  power  of  dissolved  substances  to 
absorb  light  and  the  nature  of  solution,  we  must  have  some 
quantitative  method  of  studying  the  intensities  of  the 
absorption  lines  and  bands  and  of  the  various  parts  of  the 
same  bands.  With  this  idea  in  mind  a  very  sensitive  radio- 
micrometer  was  built  and  used  to  measure  the  intensity  of 
absorption. 

The  Radiomicrometer. — The  radiomicrometer  not  only 
provides  us  with  a  method  of  studying  absorption  spectra 
quantitatively,  but  greatly  extends  the  range  of  wave- 
lengths that  can  be  studied.  The  earlier  work  with 
the  very  sensitive  radiomicrometer  had  to  do  with  the  study 
of  solutions  of  neodymium  salts.  The  effect  of  dilution 
on  absorption  spectra  was  also  investigated  quantitatively 
by  means  of  the  radiomicrometer.  It  was  found  by  this 
method,  as  with  the  grating  and  photographic  plate,  that 
the  more  concentrated  the  solution  the  broader  the  absorp- 
tion bands.  It  was  also  found  that  in  the  more  dilute 
solution,  while  the  absorption  bands  were  narrower,  they 
were  more  intense.  Further,  in  the  more  dilute  solutions 
the  centers  of  the  bands  were  displaced  towards  the  longer 
wave-lengths. 

Solutions  More  Transparent  than  Pure  Water.  —  The 
most  interesting  and  important  result  brought  out  by  the 
work  of  Jones  and  Guy  was  the  effect  of  the  dissolved  sub- 
stance on  the  absorption  of  light  by  water.  They  noted 
that  aqueous  solutions  of  certain  hydrated  salts  are  more 
transparent  than  pure  water.  This  is  obviously  a  fact 
which  called  for  careful  study.  The  absorption  of  aqueous 
solutions  of  strongly  hydrated  salts  was  compared  with  the 
absorption  of  a  layer  of  water  equal  in  depth  to  the  water 
in  the  solution.  Similar  experiments  were  carried  out  with 
salts  which  are  only  slightly  hydrated  such  as  potas- 
sium chloride  and  ammonium  chloride  and  nitrate.  It  was 
necessary  to  select  colorless  salts  which  themselves  had 
little  or  no  absorption  in  the  infra-red  where  water  absorbs. 
It  was  found,  hi  the  earlier  work,  that  the  above-named 


SOLVATE  THEORY  OF  SOLUTION  343 

compounds  had  nearly  the  same  absorption  as  water  having 
the  same  depth  as  the  water  hi  the  solution;  but  hi  subse- 
quent work  this  conclusion  must  be  modified  for  certain 
substances  near  the  bottoms  of  the  absorption  bands. 

In  terms  of  the  solvate  theory  of  solution,  we  should 
expect  the  absorption  of  the  solution  of  a  slightly  hydrated 
salt  hi  general  not  to  differ  very  greatly  from  that  of  so 
much  pure  water,  since,  when  the  solvent  is  not  combined 
with  the  dissolved  substance,  it  is  difficult  to  see  how  either 
could  affect  appreciably  the  absorbing  power  of  the  other. 

In  the  case  of  the  strongly  hydrated  salts,  very  different 
results  were  obtained.-  As  examples  of  this  class  of  sub- 
stances calcium  and  magnesium  chlorides  and  aluminium 
sulphate  were  studied.  Take  the  results  for  a  5.3  normal 
solution  of  calcium  chloride.  The  solution  is  more  trans- 
parent from  0.9  /*  to  1  p.  It  is  again  more  transparent 
from  1.05  JLC  to  1.2  /x,  being  as  much  as  25  percent  more 
transparent  than  pure  water.  For  the  longer  wave-lengths 
the  water  is  hi  general  the  more  transparent  until  1.42  /* 
is  reached,  when  both  water  and  solution  become  equally 
opaque.  Similar  results  were  obtained  with  magnesium 
.chloride. 

The  solvate  theory  seems  to  aid  us  hi  explaining  the 
facts  just  described.  Those  compounds  which  do  not  form 
hydrates,  or  which  form  only  very  simple  hydrates,  such 
as  potassium  chloride  and  the  like,  show  results  such  as 
would  be  expected.  Their  solutions  are  not  more  trans- 
parent than  so  much  pure  water.  In  general,  the  absorp- 
tion of  such  solutions  is  of  the  same  order  of  magnitude 
as  that  of  the  water  hi  which  they  are  dissolved.  We 
shall  see  that  it  came  out  hi  later  work  that  solutions  of 
only  slightly  hydrated  salts  are  more  opaque  than  pure 
water  at  the  centers  of  the  absorption  bands.  This,  how- 
ever, does  not  affect  at  all  the  conclusions  drawn  above. 
It  is  only  the  hydrated  salts  whose  solutions  are  appre- 
ciably more  transparent  than  so  much  pure  water.  How 
does  the  solvate  theory  explain  these  facts? 


344  THE  NATURE  OF  SOLUTION 

The  combined  water  seems  to  have  less  power  to  absorb 
light  than  free  water.  This  would  account  for  the  above 
facts.  The  presence  of  the  salt  seems  to  shift  the  absorp- 
tion of  the  water  towards  the  longer  wave-lengths.  Rise 
in  temperature  and  increase  hi  concentration  shift  the 
absorption  of  the  salt  towards  the  longer  wave-lengths. 
The  effect  of  rise  hi  temperature  and  increase  in  concentra- 
tion is  to  simplify  the  hydrates  existing  in  the  solution. 
Simplifying  the  resonator,  then,  shifts  the  absorption 
toward  the  red. 

The  effect  of  the  salt  on  the  absorption  of  the  water  is 
the  same  as  rise  in  temperature  and  increase  in  the  concen- 
tration of  the  solution  on  the  absorption  of  the  dissolved 
substance.  It  may  well  be  that  the  dissolved  substance 
diminishes  the  association  of  the  solvent  and  this  sim- 
plifies the  solvent  resonator.  This  may  be  true,  especially 
with  water  of  hydration,  which  is  more  directly  under  the 
influence  of  the  dissolved  substance  than  the  free  water. 

Solutions  More  Transparent  Than  Pure  Water:  Work 
of  Shaeffer  and  Paulus.  —  The  result  obtained  by  Jones 
and  Guy  was  regarded  as  of  such  importance  in  its  bearing 
on  the  solvate  theory  of  solution,  that  is  was  thought 
desirable  to  repeat  and  elaborate  with  improved  method 
the  work  which  led  to  it.  Certain  details  of  method  and 
manipulation  were  carefully  studied,  and  the  degree  of 
accuracy  of  the  procedure  adopted  was  carefully  ascer- 
tained. The  non-hydrating  or  slightly  hydrating  salts, 
potassium  chloride,  ammonium  bromide,  and  sodium 
nitrate,  were  studied.  The  strongly  hydrated  calcium 
chloride,  magnesium  chloride,  magnesium  bromide,  mag- 
nesium sulphate,  magnesium  nitrate,  zinc  sulphate,  and 
zinc  nitrate  were  investigated  at  varying  concentrations 
and  depths  of  layers. 

Solutions  of  the  strongly  hydrated  salts  have  in  general 
greater  transparency  than  pure  water,  especially  at  the 
centers  of  the  absorption  bands.  As  the  regions  of  intense 
absorption  are  approached  in  the  longer  wave-lengths, 


SOLVATE  THEORY  OF  SOLUTION  345 

the  solution  is  much  more  transparent  than  the  pure 
solvent.  This  difference  may  amount  to  as  much  as  40 
percent. 

The  non-hydrated  or  only  slightly  hydrated  salts  give 
results  which,  in  many  respects,  are  exactly  the  opposite 
of  those  obtained  with  hydrated  salts.  In  the  three  cases 
studied,  the  solution  had  greater  absorption  than  the  sol- 
vent at  the  centers  of  the  bands.  This  is  precisely  the  op- 
posite of  what  was  found  for  the  strongly  hydrated  salts. 
Regions  of  the  spectrum,  for  which  solutions  of  hydrated 
salts  were  as  much  as  40  percent  more  transparent  than 
the  solvent,  show  for  non-hydrated  salts  that  the  solution 
is  40  percent  less  transparent. 

It  was  pointed  out  that  the  results  obtained  could  be 
best  explained  by  the  solvate  theory  of  solution.  Indeed, 
this  evidence  is  of  the  very  strongest  for  that  theory.  In 
the  solutions  studied,  more  than  half  of  the  water  was 
shown  to  be  combined  with  the  dissolved  substance.  It 
was  shown  that  this  would  certainly  alter  the  vibrational 
frequency  or  resonance  of  the  absorbing  systems. 

The  transmission  curves  obtained  seem  to  justify  the 
conclusion  that  combined  water  has  less  power  to  absorb 
light  than  uncombined.  No  other  rational  explanation  has 
been  found  which  would  account  satisfactorily  for  these 
results.  The  difference  hi  the  behavior  of  hydrated  and 
non-hydrated  salts  seems  unquestionable. 

We  regard,  then,  the  spectroscopic  evidence  in  its  bear- 
ing on  the  solvate  theory  of  solution  as  very  important. 
The  presence  of  definite  "solvent  bands"  hi  the  different 
solvents  and  the  difference  between  the  absorption  of  aque- 
ous solutions  of  non-hydrated  and  strongly  hydrated  salts 
are  to  be  counted  as  among  the  strongest  and  most  direct 
lines  of  evidence  thus  far  brought  to  light  hi  this  laboratory 
bearing  on  the  solvate  theory  of  solution. 

Summary  of  the  Lines  of  Evidence  Bearing  on  the 
Solvate  Theory  of  Solution.  —  The  following  lines  of  evi- 
dence bearing  on  the  solvate  theory  of  solution  have,  then, 


346  THE  NATURE  OF  SOLUTION 

been  established1;   a  few  having  been  discussed  in  this 
volume. 

1.  Relation  between  lowering  of  the  freezing-point  of 
water  and  water  of  crystallization  of  the  dissolved  sub- 
stance. 

2.  Approximate  composition  of  the  hydrates  formed  by 
various  substances  in  solution. 

3.  Relation  between  the  minima  in  the  freezing-point 
curves  and  the  minima  in  the  boiling-point  curves. 

4.  Relation  between  water  of  crystallization  and  tem- 
perature of  crystallization. 

5.  Hydrate  theory  in  aqueous  solutions  becomes  the 
solvate  theory  in  solutions  in  general. 

6.  Temperature  coefficients  of  conductivity  and  hydra- 
tion. 

7.  Relation  between  hydration  of  the  ions  and  their 
volumes. 

8.  Hydration  of  the  ions  and  the  velocities  with  which 
they  move. 

9.  Dissociation    as    measured    by    the    freezing-point 
method  and  by  the  conductivity  method. 

10.  Effect  of  one  salt  with  hydrating  power  on  the 
hydrates  formed  by  a  second  salt  in  the  same  solution. 

11.  Investigations  in  mixed  solvents. 

12.  Spectroscopic    evidence    bearing    on    the    solvate 
theory  of  solution;   work  of  Jones  and  Uhler. 

13.  Work  of  Jones  and  Anderson  on  absorption  spectra, 
in   which   the    presence   of    " solvate"    bands   was    first 
detected.    This  showed  that  the  solvate  had  an  effect 
on  the  absorption  of  light,  and  this  could  be  explained 
only  as  due  to  a  combination  between  the  solvent  and  the 
resonator,  or  something  containing  the  resonator. 

14.  The  work  of  Jones  and  Strong  on  absorption  spectra 
established  the  existence  of  a  larger  number  of  " solvent'' 
bands.    They  showed  that  these  were  formed  by  many 
salts  and  in  many  solvents.    They  could  even  distinguish 

1  See  Journ.  Franklin  Inst.,  176  (1913). 


SOLVATE  THEORY  OF  SOLUTION  347 

between  the  bands  of  a  salt  in  a  given  alcohol  and  in  its 
isomer.  This  was  regarded  as  very  important.  The  tem- 
perature work  of  Jones  and  Strong  was  strong  evidence 
for  the  solvate  theory. 

15.  The  work  of  Jones  and  Guy  on  the  effect  of  high 
temperature  on  the  absorption  spectra  of  aqueous  solu- 
tions, and  also  on  the  effect  of   dilution,  led  to  results 
which  were  all  in  keeping  with  the  solvate  theory. 

The  most  important  spectroscopic  work  of  Jones  and 
Guy  which  bears  on  the  solvate  theory  of  solution  is  that 
in  which  the  radiomicrometer  was  used.  It  was  here 
shown  that  solutions  of  certain  strongly  hydrated  non- 
absorbing  salts  are  more  transparent  than  pure  water 
having  a  depth  equal  to  that  of  the  water  in  the  solution. 
In  the  case  of  non-hydrated  salts  the  solution  was  the 
more  opaque.  This  shows  that  water  in  combination  with 
the  dissolved  substance  —  water  of  hydration  —  has  less 
absorption  than  pure,  free  water.  This  is  regarded  as 
striking  evidence  that  some  of  the  water  in  the  presence 
of  salts  which  are  shown  by  other  methods  to  hydrate  is 
different  from  pure,  free,  uncombined  water;  and  the 
simplest  explanation  seems  to  be  that  this  is  the  combined 
water,  or  the  water  of  bydration. 

16.  The  work  of  Jones  and  Guy  was  repeated  and 
extended  by  Jones,  Shaeffer  and  Paulus.    They  obtained 
results  of  the  same  general  character  as  those  found  by 
Jones    and    Guy.    Solutions    of   hydrated    salts    were   in 
general  more  transparent  than  pure  water,  especially  at 
the  centers  of  the  absorption  bands.    Solutions  of  non- 
hydrated  or  only  slightly  hydrated  salts  are  more  opaque 
than  pure  water,  especially  at  the  centers  of  the  bands. 

The  above  sixteen  lines  of  evidence1  all  point  to  the 
general  correctness  of  the  view  that  when  a  salt  is  dis- 
solved in  a  solvent  there  is  more  or  less  combination 
between  the  salt,  or  the  ions  resulting  from  it,  and  the 

1  For  a  discussion  of  a  number  of  these,  reference  must  be  had  to  the  indi- 
vidual papers,  or  to  the  Publications  of  the  Carnegie  Institution  of  Washington. 


348  THE  NATURE  OF  SOLUTION 

solvent.  The  magnitude  of  this  solvation  depends  upon 
the  nature  of  the  substance  and  of  the  solvent. 

The  limits  of  this  volume  will  not  permit  the  discussion 
of  the  development  of  the  theory  of  hydration  in  solution 
along  other  lines  than  those  already  considered.  Refer- 
ence should  be  made,  however,  to  the  pioneer  work  of 
Pickering,1  Nernst,2  and  Noyes,3  as  well  as  to  the  more 
recent  contributions  of  Lobry  de  Bruyn,4  Bousfield,8  Buch- 
bock,6  Morgan  and  Kanolt,7  Denison  and  Steele,8  F. 
Dolezalek,9  Washburn,10  Riesenfeld,11  and  others. 

How  the  Present  Solvate  Theory  of  Solution  Differs 
from  the  Older  Hydrate  Theory.  —  The  present  solvate 
theory  of  solution  is  not  simply  one  of  several  possible 
suggestions  which  accounts  for  a  certain  class  of  experi- 
mental facts.  It  is  the  only  suggestion  that  has  thus  far 
been  made  which  seems  to  account  satisfactorily  for  all 
of  the  facts  established.  Most  of  the  above  sixteen  lines 
of  evidence  bearing  on  solvation  in  solution  were  obtained 
as  the  direct  result  of  experimental  work  suggested  by  the 
solvate  theory  and  carried  out  to  test  this  theory.  Many 
of  the  results  were  predicted  from  this  theory  before  a 
single  experiment  was  carried  out.  Solvation,  then,  being 
accepted,  as  now  seems  pretty  generally  the  case,  the 
question  arises,  how  does  the  present  solvate  theory  of 
solution  differ  from  the  older  hydrate  theory  of  Mendel£eff, 
which  has  long  since  been  abandoned  as  untenable? 

Mendel6eff  s  theory  was  that  certain  hydroscopic  sub- 

J.  Chem.  Soc.,  57,  64  and  331  (1890);  63,  141,  436,  890,  and  998  (1893). 

Nachr.  Gesells.  Wissensch.  Gottingen,  66,  86  (1900) . 

See  books  by  this  author  and  Publications  of  the  Carnegie  Institution  of 
Washington. 

Rec.  trav.  chim.,  22,  430  (1903). 

Zeit.  phys.  Chem.,  63,  257  (1905);  J.  Chem.  Soc.,  106,  600  and  1809 
(1914). 

Zeit.  phys.  Chem.,  65,  563  (1906). 

J.  Amer.  Chem.  Soc.,  28,  572  (1906). 

Zeit.  phys.  Chem.,  67,  110  (1907). 

Ibid.,  64,  727  (1908);  71,  191  (1910);  83,  45  (1913). 

10  J.  Amer.  Chem.,  Soc.  31,  (1909). 

11  Zeit.  phys.  Chem.,  66,  672  (1909);  Z.  anorg.  Chem,  85,  401  (1914). 


SOLVATE  THEORY  OF  SOLUTION  349 

stances,  such  as  calcium  chloride,  sulphuric  acid,  and  the 
like,  formed  a  few  definite  hydrates  when  hi  the  presence 
of  water.  Thus,  sulphuric  acid  formed  the  hydrates 
H2SO4.2H20,  H2S04.6H20,  H2SO4.100H2O. 

This  view  of  Mendele"eff  was  proposed,  as  we  have  seen, 
as  the  result  especially  of  measuring  the  specific  gravities 
of  aqueous  solutions  of  such  compounds  at  different  dilu- 
tions. When  the  specific  gravities  were  plotted  against 
the  concentrations,  the  curve  was  not  a  continuous  one, 
but  showed  a  number  of  breaks.  These  breaks  MendelSeff 
could  account  for  by  assuming  that  certain  definite  hy- 
drates or  compounds  between  water  and  the  dissolved 
substances  existed  at  these  concentrations.  But  before 
a  suggestion  becomes  a  theory  there  should  be  a  fair 
amount  of  evidence  supporting  it,  showing  not  only  that 
the  suggestion  accounts  for  the  facts,  but  that  it  is  the 
only  suggestion  which  will  account  for  them.  This  was 
lacking  hi  the  so-called  MendelSeff  hydrate  theory. 

The  present  solvate  theory  of  solution  may  claim  to 
have  a  fairly  good  experimental  support,  as  the  above 
review  of  some  of  the  evidence  obtained  hi  this  laboratory 
will  show.  In  aqueous  solutions  hydration  is  a  general 
phenomenon.  Some  substances  combine  with  very  little 
water,  but  most  salts  combine  with  very  large  amounts 
of  water,  the  amount  of  combined  water  for  any  given 
substance  being  a  function  of  the  concentration  of  the 
solution  and  of  the  temperature.  The  more  dilute  the 
solution,  the  larger  the  amount  of  the  solvent  combined 
with  the  dissolved  substance,  i.e.,  the  more  complex  the 
solvate.  The  lower  the  temperature  the  more  complex 
he  solvate.  These  solvates  are  very  unstable;  indeed, 
so  unstable  that  it  seems  better  to  call  them  systems  than 
definite  chemical  compounds.  Anything  so  easily  broken 
down  by  rise  in  temperature  could  hardly  be  called  a  chemi- 
cal compound.  Here,  again,  the  present  solvate  theory 
differs  from  the  older  hydrate  theory. 

While  there  is  some  spectroscopic  evidence  pointing  to 


350  THE  NATURE  OF  SOLUTION 

the  existence  in  solution  of  a  certain  definite  hydrate,  or 
certain  definite  hydrates,  we  have  obtained  a  large  amount 
of  evidence  which  seems  to  indicate  the  existence  hi  aqueous 
solutions  of  a  large  number  of  hydrates,  or  indeed  of  a  whole 
series  of  hydrates,  the  composition  depending  primarily 
on  the  concentration  of  the  solution.  While  this  is  not 
essential  to  the  present  solvate  theory  of  solution,  it  would 
differentiate  it  fundamentally  from  the  older  hydrate  theory. 

The  present  theory  is  not  simply  a  hydrate  theory  of 
aqueous  solutions.  Evidence  has  been  obtained,  and  is 
herein  briefly  discussed,  which  shows  that  solvents  other 
than  water  combine  with  the  dissolved  substance.  This 
has  been  established  for  the  alcohols  by  the  boiling-point 
method,  and  for  the  alcohols  and  many  other  solvents  by 
spectroscopic  investigations.  Indeed,  enough  evidence  has 
already  been  obtained  to  make  it  highly  probable  that 
solvation  is  not  limited  to  aqueous  solutions,  but  is  a 
general  property  of  solutions.  Solvents  in  general  have 
more  or  less  power  to  combine  with  substances  dissolved 
in  them  —  in  a  word,  we  have  the  solvate  instead  of  simply 
a  hydrate  theory. 

The  evidence  pointing  to  the  general  correctness  of 
the  solvate  theory  of  solution  is,  then,  so  strong  that  it 
seems  that  this  conception  is  in  accord  with  a  fundamental 
condition  in  connection  with  the  nature  of  solution. 

The  question  now  arises,  of  what  scientific  significance 
or  value  is  the  establishing  of  the  fact  that  there  is  more 
or  less  combination  between  the  dissolved  substance  and 
the  solvent? 

Significance  of  the  Solvate  Theory  of  Solution.1  —  The 
evidence  for  the  solvate  theory  of  solution,  which  has  been 
furnished  in  this  laboratory  as  the  result  of  somewhat 
more  than  a  dozen  years  of  investigation,  has  recently 
been  brought  together  and  briefly  discussed.2  The  evi- 

1  The  following  pages  are  taken  directly  from  the  Author's  paper  in  the 
Journ.  of  the  Franklin  Inst.,  176  (1913). 

2  Zeit.  phys.  Chem.,  74,  325  (1910). 


SOLVATE  THEORY  OF  SOLUTION  351 

dence  is  so  unambiguous  and  convincing,  that  ions  and  some 
molecules  combine  with  more  or  less  of  the  solvent,  that  it 
seems  that  it  can  now  be  accepted  as  a  fact  of  science. 

This,  however,  raises  a  number  of  questions:  What 
relation  does  the  solvate  theory  of  solution  bear  to 
the  theory  of  electrolytic  dissociation?  Does  the  solvate 
theory  help  us  to  explain  any  of  the  apparent  dis- 
crepancies in  the  theory  of  electrolytic  dissociation? 
Does  the  solvate  theory  help  us  to  explain  the  facts 
of  chemistry  hi  general  and  of  physical  chemistry  in 
particular?  Why  is  the  nature  of  solution  so  important, 
not  only  for  chemistry  but  for  science  hi  general? 
*  The  Solvate  Theory  and  the  Theory  of  Electrolytic 
Dissociation.  —  When  Arrhenius  proposed  the  theory  of 
electrolytic  dissociation,  the  question  was  not  even  raised 
as  to  the  condition  of  the  ions  hi  the  solution,  except 
that  they  behave  as  if  they  existed  independently  of  one 
another  in  solution.  The  theory  simply  said  that  mole- 
cules of  acids,  bases  and  salts,  in  the  presence  of  a  dis- 
sociating solvent  like  water,  break  down  to  a  greater  or 
less  extent  into  charged  parts  called  ions,  the  cations  or 
positively  charged  parts  being  electrically  equivalent  to 
the  anions  or  negatively  charged  parts.  The  cations  are 
usually  simple  metallic  atoms  carrying  one  or  more  unit 
charges  of  positive  electricity.  The  cation  might,  however, 
be  more  or  less  complex,  as  illustrated  by  ammonium  and 
its  substitution  products.  The  anion  is  usually  complex, 
consisting  of  a  larger  or  smaller  number  of  atoms.  It 
may,  however,  be  an  atom  carrying  negative  electricity, 
as  hi  the  case  of  the  halogen  acids  and  their  salts. 

The  degree  of  dissociation  is  determined  by  the  nature 
of  the  acid,  base  or  salt.  Strong  acids  and  bases  are 
greatly  dissociated.  Indeed,  the  degree  of  dissociation 
determines  their  strength.  Nearly  all  of  the  salts  are 
strongly  dissociated  compounds,  there  being,  however, 
some  exceptions,  as,  notably,  the  halogen  salts  of  mer- 
cury, cadmium,  and  zinc.  There  are,  however,  consider- 


352  THE  NATURE  OF  SOLUTION 

able  differences  in  the  amounts  to  which  salts  in  general 
are  dissociated  at  the  same  dilution. 

The  quantitative  evidence  furnished  by  Arrhenius  and 
others  for  the  theory  of  electrolytic  dissociation  is  so  con- 
vincing that  few  chemists  of  any  prominence,  who  have 
carefully  examined  the  evidence,  have  ever  doubted  the 
general  validity  of  the  theory;  and  the  theory  has  become 
substantiated  by  such  an  abundance  of  subsequently  dis- 
covered facts  that  it  has  now  become  a  law  of  nature  and  a 
fundamental  law  of  chemical  science. 

Arrhenius  saw  and  pointed  out  clearly  the  importance 
of  ions  for  chemistry;  Ostwald  and  his  pupils  have  shown 
that  chemistry  is  essentially  a  science  of  the  ion,  molecules 
for  the  most  part  being  incapable  of  reacting  chemically 
with  molecules;  and  Nernst  has  proved  that  the  ion  is 
the  active  agent  in  all  forms  of  primary  cells. 

The  theory  of  electrolytic  dissociation,  as  already  stated, 
does  not  raise  the  question  as  to  the  relation  between  the 
ion  and  the  solvent.  At  the  time  that  the  theory  was 
proposed,  chemists  did  not  know,  and  probably  had  no 
means  of  finding  out,  whether  the  ion  is  or  is  not  combined 
with  the  solvent  in  contact  with  it.  The  solution  of  this 
problem  remained  for  subsequent  work. 

The  solvate  theory  of  solution  has  been  regarded  in 
some  quarters  as  a  rival  of  the  electrolytic  dissociation 
theory  of  solution,  if  not  directly  antagonistic  to  it.  Such 
is  not  at  all  the  case.  The  solvate  theory  begins  where 
the  theory  of  electrolytic  dissociation  ends.  The  latter 
gives  us  the  ions  from  molecules,  and  the  former  tells  us 
the  condition  of  the  ions  in  the  presence  of  a  solvent  after 
they  are  formed. 

The  solvate  theory  of  solution,  then,  simply  supple- 
ments the  theory  of  electrolytic  dissociation,  and  both  must 
be  taken  into  account  if  we  ever  wish  to  understand  the 
phenomena  presented  by  solution. 

Does  the  Solvate  Theory  Help  to  Explain  Any  of  the 
Apparent  Exceptions  to  the  Theory  of  Electrolytic  Dis- 


SOLVATE  THEORY  OF  SOLUTION  353 

sociation?  —  Given  the  theory  of  solvation  in  solution 
together  with  that  of  electrolytic  dissociation,  the  first 
question  that  arises  is,  does  the  former  really  aid  us  in 
explaining  the  phenomena  presented  by  solutions? 

Shortly  after  the  theory  of  electrolytic  dissociation  was 
proposed,  it  was  recognized  and  repeatedly  pointed  out 
that  after  all  it  is  only  a  theory  of  "ideal  solutions,"  i.e., 
very  dilute  solutions.  It  was  shown  not  to  be  able  to 
explain  many  of  the  phenomena  presented  by  even  fairly 
concentrated  solutions.  Indeed,  it  frequently  could  not 
deal  quantitatively  with  the  very  solutions  with  which  we 
work  hi  the  laboratory.  The  explanation  of  this  short- 
coming was  not  fully  seen,  and  an  analogy  was  resorted  to. 
It  was  pointed  out  that  the  laws  of  Boyle  and  Gay-Lussac 
for  gases  hold  only  for  "ideal  gases,"  i.e.,  dilute  gases,  but 
do  not  hold  for  gases  of  any  considerable  concentration. 

It  was  stated  that  the  gas  laws  when  applied  to  solu- 
tions could  not  be  expected  to  hold  more  generally  than 
when  applied  to  gases,  and  there  the  matter  was  allowed 
to  rest. 

In  the  early  days  of  the  theory  of  electrolytic  dissocia- 
tion it  was,  however,  pointed  out  that  we  have  a  fairly 
satisfactory  explanation  of  why  the  simple  gas  laws  do  not 
hold  for  concentrated  gases,  this  being  expressed  hi  the 
equation  of  Van  der  Waals;  while  no  analogous  explana- 
tion was  offered  in  the  case  of  solutions.  That  this  point 
was  well  taken  is  obvious.  A  theory  of  solution,  to  be 
of  the  greatest  value,  must  be  applicable  to  all  solutions, 
regardless  of  the  nature  of  the  substance,  regardless  of  the 
nature  of  the  solvent,  and  regardless  of  the  concentration 
of  the  solution. 

An  explanation  of  these  apparent  exceptions  to  the 
theory  of  electrolytic  dissociation  presented  by  concen- 
trated solutions  has  been  furnished  by  the  solvate  theory. 
We  now  know  that,  for  solutions  in  general,  a  part  of  the 
solvent  is  combined  with  the  dissolved  substance.  While 
the  amount  of  the  solvent  combined  with  any  one  ion  is 


354  THE  NATURE  OF  SOLUTION 

greater  the  more  dilute  the  solution,  at  least  up  to  a  certain 
point,  the  total  amount  of  the  solvent  in  combination  with 
the  dissolved  substance  is  greater  the  more  concentrated 
the  solution. 

That  the  amount  of  combined  solvent  may  become  very 
great,  even  relative  to  the  total  amount  of  solvent  present, 
can  be  seen  from  the  folio  whig  facts:  In  a  normal  solution 
of  calcium  chloride  about  two-fifths  of  the  total  water 
present  is  combined  with  the  dissolved  substance.  In  a 
three-normal  solution  of  calcium  chloride  about  five- 
sevenths  of  the  total  water  is  combined. 

In  the  case  of  a  normal  solution  of  aluminium  chloride  in 
water,  about  five-eighths  of  the  water  present  is  combined 
with  the  dissolved  substance,  while  in  a  two-normal 
solution  about  five-sixths  of  the  water  present  is  in  a  state 
of  combination. 

What  we  suppose  to  be  a  normal  solution  of  calcium 
chloride  is,  therefore,  more  than  one  and  one-half  times 
normal,  while  what  we  suppose  to  be  a  three-normal  solu- 
tion is  in  reality  more  than  eight  times  normal.  In  the 
case  of  aluminium  chloride,  what  we  suppose  to  be  a  nor- 
mal solution  is  more  than  twice  normal,  while  what  we 
prepare  as  a  two-normal  solution  is  about  twelve  times 
normal. 

These  few  facts,  taken  from  thousands  of  a  similar 
character,  show  that  even  fairly  concentrated  solutions  are 
much  more  concentrated  than  we  would  suppose  from  the 
method  of  their  preparation;  while  very  concentrated 
solutions  are  many  times  more  concentrated  than,  without 
the  facts  of  solvation,  we  should  be  led  to  expect. 

The  general  conclusion  is  that  even  fairly  concentrated 
solutions  are  much  stronger  than  if  no  solvation  occurred, 
and  are  much  more  concentrated  than  we  are  accustomed 
to  consider  them  to  be  from  the  amount  of  substance  added 
to  a  given  volume  of  the  solvent  —  more  or  less  of  the 
water  present  being  in  combination  and  only  the  remainder 
playing  the  r61e  of  solvent.  Without  the  theory  of  solva- 


SOLVATE  THEORY  OF  SOLUTION  355 

tion,  we  have  hitherto  regarded  all  of  the  water  present  as 
acting  as  solvent. 

We  should,  therefore,  not  expect  the  laws  of  gases  to 
apply  to  such  solutions,  when  we  had  no  idea  what  was 
their  concentration.  Now  that  we  know  their  concentra- 
tion, we  may  find  that  the  laws  of  gases  are  of  as  general 
applicability  to  solutions  as  to  gases,  holding  not  simply 
for  dilute,  but  also  for  concentrated  solutions. 

The  theory  of  electrolytic  dissociation,  supplemented  by 
the  theory  of  solvation,  is,  then,  not  simply  a  theory  of  dilute 
or  "ideal"  solutions,  but  a  theory  of  solutions  in  general. 

Does  the  Solvate  Theory  Aid  in  Explaining  the  Facts 
of  Chemistry  in  General  and  of  Physical  Chemistry  in 
Particular?  —  To  answer  this  question  at  all  fully  would 
lead  us  far  beyond  the  scope  of  this  volume.  A  few 
facts  bearing  upon  this  question  can,  however,  be  taken  up. 
Take,  for  example,  the  action  of  the  hydrogen  ion  both  hi 
the  formation  and  saponification  of  esters.  In  the  pres- 
ence of  the  alcohols  the  hydrogen  ion  accelerates  greatly 
the  velocity  with  which  an  ester  is  formed,  while  in  the 
presence  of  water  it  causes  the  ester  to  break  down  into 
the  corresponding  acid  and  alcohol. 

In  terms  of  ordinary  chemical  conceptions  it  is  difficult, 
not  to  say  impossible,  to  interpret  these  reactions,  the 
hydrogen  ion  under  one  set  of  conditions  undoing  what 
under  other  conditions  it  effects.  ; 

In  terms  of  the  solvation  theory  these  reactions  admit 
of  a  very  simple  interpretation.  While  the  hydrogen  ion  is 
not  strongly  solvated,  work  hi  this  laboratory  has  shown 
that  all  ions  are  more  or  less  solvated.  In  the  presence 
of  alcohol  the  hydrogen  ion  therefore  combines  with  a 
certain  amount  of  this  solvent.  The  hydrogen  ion,  plus 
the  alcohol  combined  with  it,  unites  with  the  organic  acid, 
forming  complex  alcoholated  ions  which  then  break  down 
yielding  the  ester. 

On  the  other  hand,  the  hydrogen  ion  hi  the  presence  of 
water  combines  with  a  certain  amount  of  this  solvent.  The 


356  THE  NATURE  OF  SOLUTION 

hydrated  hydrogen  ion,  together  with  the  water  united  with 
it,  combines  with  the  ester,  forming  a  complex  hydrated 
ion,  which  then  breaks  down  into  the  corresponding  acid 
and  alcohol  setting  the  hydrogen  free  again.  For  a  fuller 
discussion  of  this  reaction  see  paper  by  E.  Emmet  Reid.1 

A  reaction  analogous  to  the  above  is  that  of  hydrogen 
ions  on  amides  in  the  presence  of  water  on  the  one  hand, 
and  alcohol  on  the  other  hand.  In  the  presence  of  water 
the  hydrated  hydrogen  ion  combines  with  the  amide,  form- 
ing a  complex  hydrated  ion  which  then  breaks  down, 
yielding  ammonia  and  acid,  the  ammonia,  of  course,  com- 
bining with  the  acid. 

In  the  presence  of  alcohol  the  alcoholated  hydrogen  ion 
combines  with  the  amide,  forming  a  complex  alcoholated 
ion,  which  then  breaks  down  into  ammonia  and  the  ester 
of  the  acid  in  question. 

Hydrogen  ions  in  a  mixture  of  water  and  alcohol,  which 
would  contain  both  hydrated  and  alcoholated  hydrogen 
ions,  give  both  reactions  simultaneously;  but,  as  Reid 
has  pointed  out,  hi  the  presence  of  an  equal  number  of 
molecules  of  water  and  alcohol,  the  tendency  of  the  hydro- 
gen ion  to  hydrate  is  greater  than  the  tendency  to  form 
alcoholates;  and  under  these  conditions  the  first  reaction 
proceeds  much  more  rapidly  than  the  second.2  A  very 
large  number  of  types  of  reactions  could  be  discussed 
illustrating  this  same  point,  i.e.,  the  value  of  the  solvate 
theory  in  interpreting  chemical  reactions. 

When  we  turn  to  physical  chemical  processes,  the  sol- 
vation of  the  ions  has  to  be  taken  into  account  at  every 
turn.  The  velocities  of  the  ions  are,  of  course,  a  function 
of  the  degree  of  their  solvation;  and  the  behavior  of  the 
ions,  both  chemically  and  physically,  is  a  function  of  their 
velocities.  The  effect  of  dilution,  and  especially  of  tem- 
perature in  reaction  velocities,  is  largely  a  question  of  the 
velocities  of  the  ions  present,  which,  in  turn,  are  a  function 
of  the  degree  of  their  solvation. 

1  Amer.  Chem.  Journ.  41,  504  (1909).  »  IUd.,  41,  509  (1909). 


SOLVATE  THEORY  OF  SOLUTION  357 

In  determining  the  actual  concentration  of  a  solution, 
the  amount  of  the  solvent  combined  with  the  ions  must  be 
taken  into  account,  as  has  already  been  pointed  out;  and 
without  knowing  the  actual  concentrations  of  solutions 
quantitative  chemistry  would  be  impossible. 

The  solvate  theory  has  thrown  a  flood  of  light  on  the 
whole  subject  of  the  conductivity  of  solutions,  or  the 
power  of  the  ions  to  carry  the  electric  current.  It  has  shown 
us  why  the  conductivity  of  lithium  salts  is  less  than  that  of 
sodium  and  potassium,  notwithstanding  the  fact  that  the 
lithium  atom  is  much  smaller  and  lighter  than  the  atom  of 
sodium  or  potassium.  We  now  know  that  the  lithium  ion 
is  much  more  hydrated  than  the  ions  of  the  other  elements, 
and  the  mass  of  the  moving  ion  is  really  much  greater 
in  the  case  of  lithium. 

When  we  come  to  the  temperature  coefficients  of  con- 
ductivity, the  solvate  theory  has  enabled  us  to  interpret 
results  which,  without  its  aid,  would  be  meaningless.  We 
now  know  why  ions  with  the  greater  hydrating  power  have 
the  larger  temperature  coefficients  of  conductivity.  We 
know  why  ions  with  the  same  hydrating  power  have 
approximately  the  same  temperature  coefficients  of  con- 
ductivity, and  why  dilute  solutions  have  larger  temperature 
coefficients  of  conductivity  than  concentrated  solutions;1 
and,  did  space  permit,  we  could  multiply  examples,  almost 
without  limit,  of  the  effect  of  the  solvate  theory  on  phys- 
ical or  general  chemistry. 

Why  the  Nature  of  Solution  is  of  Such  Vital  Importance 
not  only  for  Chemistry  but  for  Science  in  General.  —  The 
fact  is  well  recognized  that  modern  physical  or  general 
chemistry  has  reached  out  into  nearly  every  branch  of 
science,  and  has  had  an  important  influence  on  many  of 
them.  The  question  arises:  Why  is  this  the  case?  The 
answer  is  that  physical  or  general  chemistry  is  primarily  a 
science  of  solutions.  And  the  importance  of  solution  for 
science  in  general  has  already  been  pointed  out  hi  the 

1  Amer.  Chem.  Journ.,  35,  445  (1906). 


358  THE  NATURE  OF  SOLUTION 

introduction  to  this  volume.  An  examination  of  the  facts 
there  referred  to  will  show  that  the  relation  of  physical  or 
general  chemistry  to  solutions  is  the  prime  reason  why 
physical  or  general  chemistry  is  so  closely  related  to  so 
many  other  branches  of  natural  science.  This  alone  would 
show  the  importance  of  a  true  and  comprehensive  theory 
of  solutions,  not  alone  for  physical  or  general  chemistry, 
but  for  the  natural  sciences  in  general. 

An  examination  of  the  literature  on  the  nature  of  solu- 
tion (Chapter  II)  has  brought  out  a  fact  of  general  sig- 
nificance hi  the  history  of  the  development  of  any  branch 
of  science,  and  of  any  subdivision  of  any  branch.  The 
evolution  of  the  subject  is  slow  and  takes  place  by  a  series 
of  small  increments,  one  suggestion  contributing  a  little 
here  and  another  a  little  there.  Sometimes  the  steps  are 
relatively  large,  but  on  the  whole  it  is  given  to  any  one 
man  of  science  to  discover  only  a  small  amount  of  new 
truth. 


BIBLIOGRAPHY  OF  ARTICLES  AND  BOOKS 

BY  PROFESSOR  H.  C.  JONES  AND  His  COWOBKERS 

MOUSE,  H.  N.,  and  JONES,  H.  C. 

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on  the  Sub-hydroxide  and  Sub-oxide  of  Cadmium.    Amer.  Chem.  Journ., 
12,  488  (1890). 
MORSE,  H.  N.,  and  JONES,  H.  C. 

A  Redetermination  of  the  Atomic  Weight  of  Cadmium.    Amer.  Chem.  Journ., 

14,  261  (1892). 
JONES,  H.  C. 

The  Uses  of  Hydrogen  Dioxide  in  Quantitative  Analysis,  and  the  Important 
Methods  for  Determining  Hydrogen  Dioxide.     Amer.  Chem.  Journ.,  IS, 
275  (1890). 
JONES,  H.  C. 

Uber  den  Gefrierpunkt  sehr  verdunnter  Losungen.     Zeit.  phys.  Chem.,  11, 

110  (1892). 
JONES,  H.  C. 

Uber  die  Bestimmung  des  Gefrierpunktea   sehr  verdunnter   Salzlosungen. 

Zeit.  phys.  Chem.,  11,  529  (1892). 
JONES,  H.  C. 

Uber  die  Bestimmung  des  Gefrierpunktes  «ehr  verdunnter  Losungen  Sauren, 
Alkalien,   Salze  und  organischen  Verbindungen.     Zeit.   phys.   Chem.,   IS, 
623  (1893). 
JONES,  H.  C. 

Zur  Bestimmung  des  Gefrierpunktea  sehr  verdunnter  Salzlosungen.     Ber.  d. 

chem.  Ge&ett.,  26,  547  (1893). 
JONES,  H.  C. 

Uber   den    Gefrierpunkt  verdunnter   Losungen  von  Chlornatrium.    Ber  d. 

chem.  Gesett.,  26,  1633  (1893). 
JONES,  H.  C. 

The  Dissociation  of    Compounds  in  Water  as  Measured  by  the   Lowering 
of  the  Freezing-point,  and  some  Investigations  with  Organic  Compounds. 
Phil.  Mag.,  36,  465  (1893). 
JONES,  H.  C. 

Die  Emiedrigung  des  Gefrierpunktes  des  Losungsmittels  durch  Electrolyte. 

Wied.  Ann.,  53,  392  (1893). 
JONES,  H.  C. 

Uber   die   Verbindung   von   Schwefelsaure   mit   Wasser   in   Gegenwart  von 

Essigsaure.     Zeit.  phys.  Chem.,  13,  419  (1893). 
JONES,  H.  C. 

On  the  Combination  of  Sulphuric  Acid  with  Water  in  the  Presence  of  Acetic 

Acid.     Amer.  Chem.  Joum.,  16,  1  (1894). 
JONES,  H.  C. 

Uber  die  Losungstension  von  Metallen.     Zeit.  Phys.  Chem.,  14,  346  (1894). 
JONES,  H.  C. 

On  the  Solution-tension  of  Metals.     Phys.  Rev.,  2,  8  (1895). 


360  THE  NATURE  OF  SOLUTION 

JONES,  H.  C. 

tiber   die   Gefrierpunktserniedrigung   verdiinnter   wasseriger   Losungen  von 

Nichtelektrolyten.     Zeit.  phys.  Chem.,  18,  283  (1895). 
JONES,  H.  C. 

A  Redetermination  of  the  Atomic  Weight  of  Yttrium.     Amer.  Chem.  Journ. 

17,  154  (1895). 
JONES,  H.  C. 

The  Determination  of  Formic  Acid  by  Titration  with  Potassium  Perman- 
ganate.    Amer.  Chem.  Journ.,  17,  539  (1895). 
JONES,  H.  C. 

The  Dilution  Law.     Amer.  Chem.  Journ.,  18,  343  (1896). 
JONES,  H.  C.,  and  ALLEN,  CHARLES  R. 

The  Use  of  Phenolphthalein  in  Illustrating  the  Dissociating  Action  of  Water. 

Amer.  Chem.  Journ.,  18,  377  (1896). 
JONES,  H.  C.,  and  MACKAY,  E. 

A  Contribution  to  the  Study  of  Water  Solutions  of  Some  of    the  Alums. 

Amer.  Chem.  Journ.,  19,  83  (1897). 
JONES,  H.  C. 

A  Simple  and  Efficient  Boiling-point  Apparatus  for  use  with  Low  and  with 

High-boiling  Solvents.     Amer.  Chem.  Journ.,  19,  581  (1897). 
JONES,  H.  C.,  and  KING,  S.  H. 

The  Dissociation  of  Electrolytes  as  Measured  by  the  Boiling-point  Method. 

Amer.  Chem.  Journ.,  19,  753  (1897). 
JONES,  H.  C. 

A  Determination  of  the  Atomic  Weight  of  Praseodymium  and  of  Neodymium. 

Amer.  Chem.  Joum.,  20,  345  (1898). 
JONES,  H.  C. 

The  Opening  of  the  New  Laboratory  for  Physical  Chemistry  in  Leipzig. 

Science,  7,  786  (1898). 
JONES,  H.  C. 

The  Rise  of  the  Theory  of  Electrolytic  Dissociation  and  a  Few  of  its  Applica- 
tions in  Chemistry,  Physics  and  Biology.     Johns  Hopkins  Hospital  Bulletin, 
June,  1898. 
JONES,  H.  C. 

Notiz  iiber  das  Atomgewicht  von  Praseodym  und  Neodym.     Zeit.  anorg. 

Chem.,  19,  339  (1899). 
JONES,  H.  C.,  and  OTA,  K. 

Contributions  to  our  Knowledge   of  Aqueous  Solutions    of   Double   Salts. 

Amer.  Chem.  Journ.,  22,  5  (1899). 
JONES,  H.  C.,  and  KNIGHT,  N. 

Contribution  to  the  Study  of  Aqueous  Solutions  of   Double  Salts.     Amer. 

Chem.  Joum.,  22,  110  (1899). 
JONES,  H.  C. 

The  Electrolytic  Dissociation  of  Certain  Salts  in  Methyl  Alcohol  as  Measured 

by  the  Boiling-point  Method.     Zeit.  phys.  Chem.,  31,  114  (1900). 
JONES,  H.  C.,  and  CHAMBERS,  V.  J. 

On  some  Abnormal  Freezing-point  Lowerings  produced    by  Chlorides  and 

Bromides  of  the  Alkaline  Earths.     Amer.  Chem.  Journ.,  23,  89  (1900). 
JONES,  H.  C.,  and  SMITH,  A.  W. 

The  Solution-tension  of  Zinc  in  Ethyl  Alcohol.     Amer.  Chem.  Journ.,  23, 
397  (1900). 


BIBLIOGRAPHY  OF  ARTICLES  AND  BOOKS  361 

CHAMBERS,  V.  J.,  and  FRAZER,  J.  C.  W. 

On  a  Minimum  in  the  Molecular  Lowering  of  the  Freezing-point  of  Water, 
produced  by  certain  Acids  and  Salts.     Amer.  Chem.  Journ.,  23,  512  (1900). 
LINDSAY,  C.  F. 

The  Conductivities  of  some  Double  Salts  as  Compared  with  the  Conduc- 
tivities of  Mixtures  of  their  Constituents.     Amer.  Chem.  Journ.,  25,  62 
(1901). 
JONES,  H.  C. 

The  Dissociating  Power  of  Different  Solvents.    Amer.  Chem.  Journ.,  25, 

232  (1901). 
JONES,  H.  C.,  and  CALDWELL,  B.  P. 

Contribution  to  the  Study  of  Aqueous  Solutions  of  Double  Salts.    Amer. 

Chem.  Journ.,  25,  349  (1901). 
JONES,  H.  C.,  and  DOUGLAS,  J.  M. 

The  Dissociation  of  Certain  Acids,  Bases  and  Salts  at  Different  Temperatures. 

Amer.  Chem.  Journ.,  26,  428  (1901). 
JONES,  H.  C.,  and  MATHER,  W.  T. 

A  New  Apparatus  for  Determining  the  Relative  Velocities  of   Ions;    with 

some  Results  for  Silver  Ions.     Amer.  Chem.  Journ.,  26,  473  (1901). 
JONES,  H.  C. 

The  Molecular  Weights  of  Certain  Salts  in  Acetone.     Amer.  Chem.  Journ., 

27,  16  (1902). 

JONES,  H.  C.,  BARNES,  J.,  and  HYDE,  E.  P. 

The  Lowering  of  the  Freezing-point  of  Aqueous  Hydrogen  Dioxide.     Amer. 

Chem.  Journ.,  27,  22  (1901). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

The  Lowering  of  the  Freezing-point  of  Water  produced  by  Concentrated 
Solutions  of  Certain  Electrolytes,  and  the  Conductivity  of  such  Solutions. 
Amer.  Chem.  Journ.,  27,  433  (1902). 
JONES,  H.  C. 

A  Redetermination  of  the  Atomic    Weight  of  Lanthanum.     Amer.   Chem. 

Journ.,  28,  23  (1902). 
JONES,  H.  C.,  and  CARROLL,  C.  G. 

The  Lowering  of  the  Freezing-point  of  Aqueous  Hydrogen  Dioxide  produced 

by  Certain  Salts  and  Acids.     Amer.  Chem.  Journ.,  28,  284  (1902). 
JONES,  H.  C.,  and  LINDSAY,  C.  F. 

A  Study  of  the  Conductivity  of  Certain  Salts  in  Water,  Methyl,  Ethyl  and 
Propyl  Alcohols,  and  in  Mixtures  of  These  Solvents.  Amer.  Chem.  Journ., 

28,  329  (1902). 
JONES,  H.  C. 

Das  Atomgewicht  des  Lanthans.     ZeU.  anorg.  Chem.,  36,  92  (1903). 
JONES,  H.  C.,  and  MURRAY,  A.  G. 

The  Association  of  a  Liquid  Diminished  by  the  Presence  of  Another  Asso- 
ciated Liquid.     Amer.  Chem.  Journ.,  30,  193  (1903). 
JONES,  H.  C.,  and  MURRAY,  A.  G. 

The  Lowering  of  the  Freezing-point  of  Aqueous  Hydrogen  Dioxide  by  Sul- 
phuric and  Acetic  Acids.     Amer.  Chem.  Journ.,  30,  205  (1903). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

The  Molecular  Lowering  of  the  Freezing-point  of  Water  produced  by  Con- 
centrated Solutions  of  Certain  Electrolytes.  ZeU.  phys.  Chem.,  46,  245 
(1903). 


362  THE  NATURE  OF  SOLUTION 

JONES,  H.  C.,  and  GETMAN,  F.  H. 

A  Study  of  the  Molecular  Lowering  of  the  Freezing-point  of  Water  produced 

by  Concentrated  Solutions  of  Electrolytes.     Phys.  Rev.,  18,  146  (1904). 
JONES,  H.  C. 

The  Effect  of  One  Associated  Solvent  on  the  Association  of  Another  Asso- 
ciated Solvent.     Boltzmann  Jubelband,  105  (1904). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

On  the  Nature  of  Concentrated  Solutions  of    Electrolytes  —  Hydrates  in 

Solution.    Amer.  Chem.  Journ.,  31,  303  (1904). 
JONES,  H.  C. 

The  Significance  of  the  Maximum  of  the  Conductivity  Curves  of  Kraus  at 

High  Temperatures.     Amer.  Chem.  Journ.,  31,  584  (1904). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

tJber  die  Existenz  von  Hydraten  in  concentrierten  wasserigen  Losungen  der 
Elektrolyte  und  einiger  Nichtelektrolyte.     Ber.  d.  chem.  Gesell,  37,  1511 
(1904). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

The  Existence  of  Hydrates  in  Solutions  of  Certain  Nonelectrolytes  and  the 
Nonexistence  of  Hydrates  in  Solutions  of  Organic  Acids.      Amer.  Chem. 
Journ.,  32,  308  (1904). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

The  Existence  of  Alcoholates  in  Solutions  of  Certain  Electrolytes  in  Alcohol. 

Amer.  Chem.  Journ.,  32,  338  (1904). 
JONES,  H.  C.,  and  GETMAN,  F.  H. 

tJber  das  Vorhandensein  von  Hydraten  in  konzentrierten  wasserigen  Ldsungen 

von  Elektrolyten.     Zeit.  phys.  Chem.,  49,  385  (1904). 
JONES,  H.  C.,  and  BASSETT,  H.  P. 

Determination  of  the  Relative  Velocities  of  the  Ions  of  Silver  Nitrate  in 
Mixtures  of  the  Alcohols  and  Water,  and  on  the  Conductivity  of  Such 
Mixtures.     Amer.  Chem.  Journ.,  32,  409  (1904). 
JONES,  H.  C.,  and  CARROLL,  C.  G. 

A  Study  of  the  Conductivities  of  Certain  Electrolytes  in  Water,  Methyl  and 
Ethyl  Alcohols,  and  Mixtures  of  These  Solvents  —  Relation  between  Con- 
ductivity and  Viscosity.     Amer.  Chem.  Journ.,  32,  521  (1904). 
JONES,  H.  C.,  and  BASSETT,  H.  P. 

The  Approximate  Composition  of  the  Hydrates  formed  by  Certain  Elec- 
trolytes in  Aqueous  Solutions  at  Different  Concentrations.     Amer.  Chem. 
Joum.,  33,  534  (1905). 
JONES,  H.  C.,  and  BASSETT,  H.  P. 

Der  Einfluss  der  Temperatur  auf  die  Kristallwassermenge  als  Beweis  fur  die 
Theorie  von  den  Hydraten  in  LSsung.     Zeit.  phys.  Chem.,  52,  231  (1905). 
JONES.  H.  C. 

A  Correction.     Phil.  Mag.,  10,  157  (1905). 
JONES,  H.  C. 

L'Existence  d'Hydrates  dans  lea  Solutioni  Aqueusea  d' Electrolytes.    Journ. 

Chim.  Phys.,  3,  455  (1905). 
JONES,  H.  C.,  and  BASSETT,  H.  P. 

The  Approximate  Composition  of  the  Hydrates  Formed  by  a  Number  of 
Electrolytes  in  Aqueous  Solutions;  together  with  a  Brief  General  Dis- 
cussion of  the  Results  thus  far  Obtained.  Amer.  Chem.  Journ.,  34,  290 
(1905). 


BIBLIOGRAPHY  OF  ARTICLES  AND  BOOKS  363 

JONES,  H.  C.,  and  WEST,  A.  P. 

A  Study  of  the  Temperature  Coefficient*  of  Conductivity  in  Aqueous  Solu- 
tions, and  on  the  effect  of  Temperature  on  Dissociation.     Amer.  Chem. 
Journ.,  34,  357  (1905). 
JONBS,  H.  C. 

The  Atomic  Weight  of  Radium  and  the  Periodic  System.    Amer.  Chem. 

Journ.,  34,  467  (1905). 
JONES,  H.  C.,  and  BINGHAM,  E.  C. 

The  Conductivity  and  Viscosity  of  Solutions  of  Certain  Salts  in  Mixtures  of 
Acetone  with  Methyl  Alcohol,  with  Ethyl  Alcohol  and  Water.     Amer. 
Chem.  Journ.,  34,  481  (1905). 
JONES,  H.  C.,  and^McMASTEB,  L. 

On  the  Formation  of  Alcoholates  by  Certain  Salts  in  Solution  in  Methyl 

and  Ethyl  Alcohols.     Amer.  Chem.  Journ.,  35,  316  (1906). 
JONES,  H.  C. 

The  Bearing  of  Hydrates  on  the  Temperature  Coefficients  of  Conductivity  of 

Aqueous  Solutions.     Amer.  Chem.  Journ.,  30,  445  (1906). 
JONES,  H.  C. 

Die  annahernde  Zusammensetzung  der  Hydrate,  welche  von  verschiedenen 
Elektrolyten  in  wasseriger  Losung  gebildet  werden.     ZeU.  phys.   Chem., 
55,  385  (1906). 
JONES,  H.  C.,  LINDSAY,  C.  F.,  and  CARHOLL,  C.  G. 

tlber  die  Leitfahigkeit  gewisser  Salze  in  Gemischten  Ldsungsmitteln;  Wasser, 

Methyl-,  Athyl-  und  Propylalkohol.     Zeit.  phys.  Chem.,  56,  129  (1906). 
JONES,  H.  C.,  and  MCMASTER,  L. 

The  Conductivity  and  Viscosity  of  Solutions  of  Certain  Salts  in  Water,  Methyl 
Alcohol,  Ethyl  Alcohol,  Acetone  and  Binary  Mixtures  of  These  Solvents. 
Amer.  Chem.  Journ.,  36,  325  (1906). 
JONES,  H.  C.,  and  ROUILLEB,  C.  A. 

The  Relative  Migration  Velocities  of  the  Ions  of  Silver  Nitrate  in  Water, 
Methyl  Alcohol,  Ethyl  Alcohol  and  Acetone,  and  in  Binary  Mixtures  of 
These  Solvents,  together  with  the  Conductivity  of  Such  Solutions.     Amer. 
Chem.  Journ.,  36,  427  (1906). 
JONES,  H.  C. 

Der  Arbeitsanteil  der  Herrn  W.  Biltz  an  der  gegenwartigen  Hydrattheorie. 

Zeit.  phys.  Chem.,  57,  244  (1906). 
JONES,  H.  C.,  BINGHAM,  E.  C.,  and  MCMASTEB,  L. 

Uber  Leitfahigkeit  und  innere  Reibung  von  Losungen  gewisser  Salze  in  den 
Losungsmittelgemischen:      Wasser,     Methylalkohol,     Athylalkohol,     und 
Aceton.     Zeit.  phys.  Chem.,  57,  193,  257  (1906). 
JONES,  H.  C.,  and  UHLEB,  H.  S. 

The  Absorption  Spectra  of  Certain  Salts  in  Aqueous  Solution  as  affected  by 
the  Presence  of  Certain  Other  Salts  with  Large  Hydrating  Power.     Amer. 
Chem.  Journ.,  37,  126  (1907). 
JONES,  H.  C.,  and  UHLEB,  H.  S. 

The  Absorption  Spectra  of  Certain  Salts  in  Nonaqueous  Solvents,  as  affected 

by  the  Addition  of  Water.     Amer.  Chem.  Journ.,  37,  244  (1907). 
JONES,  H.  C.,  and  VEAZEY,  W.  R. 

A  Possible  Explanation  of  the  Increase  in  Viscosity  which  Results  when  the 
Alcohols  are  Mixed  with  Water;  and  of  the  Negative  Viscosity  Coeffi- 
cients of  Certain  Salts  when  Dissolved  in  Water.  Amer.  Chem.  Journ., 
37,  405  (1907). 


364  THE  NATURE  OF   SOLUTION 

JONES,  H.  C.,  and  PEARCE,  J.  N. 

Dissociation  as  Measured  by  Freezing-point  Lowering  and  by  Conductivity  — 
Bearing  on  the  Hydrate  theory  —  The  Approximate  Composition  of  the 
Hydrates  formed  by  a  Number  of  Electrolytes.     Amer.  Chem.  Journ.,  38, 
683  (1907). 
JONES,  H.  C.,  and  STINE,  C.  M. 

The  Effect  of  One  Salt  on  the  Hydrating  Power  of  Another  Salt  Present  in 

the  Same  Solution.     Amer.  Chem.  Journ.,  39,  313  (1908). 
JONES,  H.  C.,  and  VEAZEY,  W.  R. 

Die  Leitfahigkeit  und  innere  Reibung  von  Losungen  gewisser  Salze  in  Wasser, 
Methylalkohol,    Athylalkohol,    Aceton,    und    binaren    Gemischen    dieser 
Losungsmittel.     Zeit.  phys.  Chem.,  61,  641  (1908). 
JONES,  H.  C.,  and  VBAZBT,  W.  R. 

Die  Leitfahigkeit  und   innere  Reibung  von   Tetraathylammoniumjodid  in 
Wasser,  Methylalkohol,  Athylalkohol,  Nitrobenzol,  und  binaren  Gemischen 
dieser  Losungsmittel.     Zeit.  phya.  Chem.,  62,  44  (1908). 
JONES,  H.  C.,  and  ANDERSON,  J.  A. 

The  Absorption  Spectra  of  Neodymium  Chloride  and  Praseodymium  Chloride 
in  Water,  Methyl  Alcohol,  Ethyl  Alcohol,  and  Mixtures  of  These  Solvents. 
Proc.  Amer.  PhUosoph.  Soc.,  47,  276  (1908). 
JONES,  H.  C.,  and  JACOBSON,  C.  A. 

The  Conductivity  and  lonization  of  Electrolytes  in  Aqueous  Solutions  as 
Conditioned   by  Temperature,    Dilution   and   Hydrolysis.     Amer.   Chem. 
Journ.,  40,  355  (1908). 
TURNER,  B.  B. 

The  Limiting  Conductivity  and  Degree  of  lonization  of  Alcoholic  Solutions. 

Amer.  Chem.  Journ.,  40,  558  (1908). 
JONES,  H.  C. 

The  Present  Status  of  the  Solvate  Theory.    Amer.  Chem.  Journ.,  41,  19  (1909.) 
JONES,  H.  C.,  and  ANDERSON,  J.  A. 

The  Absorption  Spectra  of  Solutions  of  a  Number  of  Salts  in  Water,  in  Cer- 
tain Nonaqueous  Solvents,  and  in  Mixtures  of  These  Solvents  with  Water. 
Amer.  Chem.  Journ.,  41,  163  (1909). 
JONES,  H.  C.,  and  MAHIN,  E.  G. 

The  Conductivity  of  Solutions  of  Lithium  Nitrate  in  Ternary  Mixtures  of 
Acetone,   Methyl  Alcohol,  Ethyl  Alcohol  and  Water,  together  with  the 
Viscosity  and  Fluidity  of  These  Mixtures.     Amer.  Chem.  Journ.,  41,  433 
(1909). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

The  Absorption  Spectra  of  Various  Potassium,  Uranyl,  Uranous,  and  Neo- 
dymium Salts  in  Solution,  and  the  Effect  of  Temperature  on  the  Absorp- 
tion Spectra  of  Certain  Colored  Salts  in  Solution.     Proc.  Amer.  Philosoph. 
Soc.,  48,  194  (1909). 
Jones,  H.  C.,  and  MAHIN,  E.  G. 

Conductivity  and  Viscosity  of  Dilute  Solutions  of  Lithium  Nitrate  and  Cad- 
mium Iodide  in  Binary  and  Ternary  Mixtures  of  Acetone  with  Methyl 
Alcohol,  Ethyl  Alcohol  and  Water.     Zeit.  phya.  Chem.,  69,  389  (1909). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

Die    Absorptionsspektren    gewisser    Salzlosungen.     Phys.    Zeit.,    10,    499 
(1909). 
JONES,  H.  C.,  and  SCHMIDT,  M.  R. 

Conductivity  and  Viscosity  in  Mixed  Solvents  containing  Glycerol.     Amer. 
Chem.  Journ.,  42,  37  (1909). 


BIBLIOGRAPHY  OF  ARTICLES  AND  BOOKS  365 

JONES,  H.  C.,  and  STRONG,  W.  W. 

The  Absorption  Spectra  of  Various  Salts  in  Solution,  and  the  Effect  of 

Temperature  on  such  Spectra.     Amer.  Chem.  Journ.,  43,  37  (1910). 
CLOVBB,  A.  M.,  and  JONES,  H.  C. 

The  Conductivities,   Dissociations,   and  Temperature  Coefficients  of  Con- 
ductivity between  35°  and  80°  of  Solutions  of  a  Number  of  Salts  and 
Organic  Acids.     Amer.  Chem.  Journ.,  43,  187  (1910). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

The  Absorption  Spectra  of  Solutions;   A  Possible  Method  for  Detecting  the 
Presence    of    Intermediate    Compounds   in    Chemical    Reactions.     Amer. 
Chem.  Journ.,  43,  224  (1910). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

The  Absorption  Spectra  of  Certain  Uranous  and  Uranyl  Compounds.    Phil. 

Mag.,  19,  566  (1910). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

Spectres    d'Absorption    des    Solutions.    Possibility    d'une    M&thode    pour 
Determiner  la  Presence  de  Composes  IntermMiaires  dans  les  Reactions 
Chimiques.     Journ.  Chim.  Phys.,  8,  131  (1910). 
WHITE,  G.  F.»  and  JONES,  H.  C. 

The  Conductivity  and  Dissociation  of  Organic  Acids  in  Aqueous  Solution 

at  Different  Temperatures.    Amer.  Chem.  Journ.,  44,  159  (1910). 
JONES,  H.  C. 

Im  hiesigen  Laboratorium  wahrend  der  vergangenen  zwolf  Jahre  erhaltene 
Anhaltspunkte   fur   die   Existenz   von   Solvaten    in   Losung    (Dreizehnte 
Mitteilung).    Zeit.  phys.  Chem.,  74,  325  (1910). 
WEST,  A.  P.,  and  JONES,  H.  C. 

The  Conductivity,  Dissociation  and  Temperature  Coefficients  of  Conduc- 
tivity at  35°,  50°  and  65°  of  Aqueous   Solutions   of   a  Number  of  Salts. 
Amer.  Chem.  Journ.,  44,  508  (1910). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

The  Absorption  Spectra  of  Certain  Salts  of  Cobalt,  Erbium,  Neodymium 
and  Uranium,  as  Affected  by  Temperature  and  by  Chemical  Reagents. 
Amer.  Chem.  Journ.,  45,  1,  113  (1911). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

Selective  Oxidation.     Amer.  Chem.  Journ.,  45,  36  (1911). 
KREIDER,  H.  R.  and  JONES,  H.  C. 

The  Dissociation  of  Electrolytes  in  Non-Aqueous  Solvents  as  Determined 
by  the  Conductivity  and  Boiling-Point  Methods.      Amer.  Chem.  Journ., 
45,  282  (1911). 
JONES,  H.  C. 

Sur  la  Position  de  la  Theorie  des  Solvates.    Journ.  Chim.  Phys.,  9,  217  (1911). 
WINSTON,  L.  G. 

Electrical  Induction  in  Chemical  Reactions.    Amer.  Chem.  Journ.,  45,  647 

(1911). 
WIGHTMAN,  E.  P.;  and  JONES,  H.  C. 

A  Study  of  the  Conductivity  and  Dissociation  of  Organic  Acids  in  Aqueous 
Solution  between  Zero  and  Thirty-five  Degrees.     Amer.  Chem.  Journ.,  46, 
56  (1911). 
GUY,  J.  S.,  and  JONES,  H.  C. 

Conductivity  and  Viscosity  in  Mixed  Solvents  Containing  Glycerol.     Amer. 
Chem.  Journ.,  46,  131  (1911). 


366  THE  NATURE  OF  SOLUTION 

HOSFOBD,  H.  H.,  and  JONES,  H.  C. 

The   Conductivities,    Temperature   Coefficients   of   Conductivity   and   Dis- 
sociation of  Certain  Electrolytes.     Amer.  Chem.  Journ.,  46,  240  (1911). 
WINSTON,  L.  G.,  and  JONES,  H.  C. 

The  Conductivity,  Temperature  Coefficients  of  Conductivity  and  Dissocia- 
tion of  Certain  Electrolytes  in  Aqueous  Solution  from  0°  to  35°.     Prob- 
able Inductive  Action  in  Solution,  and  Evidence  for  the  Complexity  of 
the  Ion.     Amer.  Chem.  Journ.,  46,  368  (1911). 
KBEIDBB,  H.  R.,  and  JONES,  H.  C. 

The  Conductivity  of  Certain  Salts  in  Methyl  and  Ethyl  Alcohols  at  High 

Dilutions.     Amer.  Chem.  Journ.,  46,  574  (1911). 
JONES,  H.  C. 

The  Introduction  of  Physical  Chemical  Conceptions  in  the  Early  Stages  of 

the  Teaching  of  Chemistry.     Science,  35,  87  (1912). 
JONES,  H.  C.,  and  STRONG,  W.  W. 

The  Absorption  Spectra  of  Comparatively  Rare  Salts.     The  Spectropho- 
tography  of  Certain  Chemical  Reactions,  and  the  Effect  of  High  Tem- 
perature on   the  Absorption   Spectra  of  Nonaqueous  Solutions.     Amer. 
Chem.  Journ.,  47,  27,  126  (1912). 
JONES,  H.  C. 

The  Nature  of  Solution.    Journal  Franklin  Institute,  173,  217  (1912). 
JONES,  H.  C. 

Absorption  Spectra  and  the  Solvate  Theory  of  Solution.    Phil.  Mag.,  23, 

730  (1912). 
JONES,  H.  C. 

Die  Absorptionsspektra  von  Losungen.     Zeit.  Phya.  Chem.,  80,  361  (1912). 
WIGHTMAN,  E.  P.,  and  JONES,  H.  C. 

A  Study  of  the  Conductivity  and  Dissociation  of  Certain  Organic  Acids  at 

35°,  50°  and  65°.     Amer.  Chem.  Journ.,  48,  320  (1912). 
JONES,  H.  C.,  and  GUT,  J.  S. 

Die  Absorptionsspektren  wassriger  Losungen  von  Neodym-  und  Praseodym- 

salzen,  mit  dem  Radiomikrometer  gemessen.    Phys.  Zeit.,  13,  649  (1912). 
SPBINGEB,  A.,  JR.,  and  JONES,  H.  C. 

A  Study  of  the  Conductivity  and   Dissociation  of  Certain  Organic  Acids  in 
Aqueous  Solution  at  Different  Temperatures.     Amer.  Chem.  Journ.,  48, 
411  (1912). 
HOWABD,  S.  F.,  and  JONES,  H.  C. 

The  Conductivity,  Temperature  Coefficients  of  Conductivity  and  Dissocia- 
tion of  Certain  Electrolytes  in  Aqueous  Solution  at  35°,  50°  and  65°. 
Amer.  Chem.  Journ.,  48,  500  (1912). 
DAVIS,  P.  B.,  and  JONES,  H.  C. 

Leitfahigkeits-  und  negative  Viskositatskoeffizienten  gewisser  Rubidium-  und 
Ammoniumsalze  in  Glycerin  und  in  Gemischen  von  Glycerin  mit  Wasser 
25°  bis  75°.     Zeit.  phys.  Chem.,  81,  68  (1913). 
JONES,  H.  C.,  and  GUY,  J.  S. 

The  Absorption  Spectra  of  Solutions  as  Affected  by  Temperature  and  by 
Dilution.     A  Quantitative  Study  of  Absorption  Spectra  by  Means  of  the 
Radiomicrometer.     Amer.  Chem.  Journ.,  49,  1  (1913). 
SHAEFFEB,  E.  J.,  and  JONES,  H.  C. 

A  Study  of  the  Conductivity,  Dissociation  and  Temperature  Coefficients 
of  Conductivity  of  Certain  Inorganic  Salts  in  Aqueous  Solution,  as  Con- 
ditioned by  Temperature,  Dilution,  Hydration  and  Hydrolysis.  Amer. 
Chem.  Journ.,  49,  207  (1913). 


BIBLIOGRAPHY  OF  ARTICLES   AND    BOOKS  367 

JONES,  H.  C. 

The  Bearing  of  Osmotic  Pressure  on  the  Development  of  Physical  or  General 

Chemistry.     The  Plant  World,  16,  73  (1913). 
JONES,  H.  C. 

L'Influence  de  la  Pression  Osmotique  sur  le  Development  de  la  Chimie 

Physique.     Revue  Scientifaue,  Oct.  11  (1913). 
GUT,  J.  S.,  SHAEFTER,  E  J.,  and  JONES,  H.  C. 

The  Absorption  of  Light  by  Water  Changed  by  the  Presence  of  Strongly 
Hydrated  Salts,  as  Shown  by  the  Radiomicrometer  —  New  Evidence  for 
the  Solvate  Theory  of  Solution.     Amer.  Chem.  Journ.,  49,  265  (1913). 
GUT,  J.  S.,  SHAEFFEB,  E.  J.,  and  JONES,  H.  C. 

Die  Anderung  der  Absorption  des  Lichtes  durch  Wasser  infolge  der  Gegenwart 
stark  hydrieter  Salze,  nachgewiesen  mit  Hilfe  des  Radiomikrometers  —  Ein 
neuer  Bewies  fur  die  Solvattheorie  der  Losungen.     Phys.  Zeit.,  14,  278 
(1913). 
SMITH,  L.  D.,  and  JONES,  H.  C. 

Conductivity,  Temperature  Coefficients  of  Conductivity,   Dissociation  and 
Dissociation  Constants  of  Certain  Organic  Acids  between  0°  and  65°. 
Amer.  Chem.  Journ.,  60,  1  (1913). 
GUT,  J.  S.,  and  JONES,  H.  C. 

A  Quantitative  Study  of  Absorption  Spectra  by  Means  of  the  Radiomicro- 
meter.    Amer.  Chem.  Journ.,  50,  257  (1913). 
JONES,  H.  C. 

Evidence  Bearing  on  the  Solvate  Theory  of  Solution.     Journal  Franklin 

Institute,  176,  479,  677  (1913). 
DAVIS,  P.  B.,  HUGHES,  H.,  and  JONES,  H.  C. 

Leitfahigkeit  und  Viskositat  von  Losungen  von  Rubidiumsalzen  in  Gemischen 

von  Aceton  und  Wasser.     Zeit.  phys.  Chem.,  85,  513  (1913). 
SCHAEFFER,  E.  JM  PAULUS,  M.  G.,  and  JONES,  H.  C. 

Die  Anderung  der  Absorption  des  Lichtes  durch  Wasser  infolge  der  Gegen- 
wart stark  hydrieter  Salze,   gemessen  mit  Hilfe  des  Radiomikrometers. 
Phys.  Zeit.,  15,  447  (1914). 
JONES,  H.  C.,  and  GUT,  J.  S. 

Eine    quantitative    Untersuchung    der   Absorptionsspektren    von    Losungen 

mittels  des  Radiomikrometers.     Ann.  d.  Phys.,  43,  555  (1914). 
JONES,  H.  C. 

Absorptionsspektra  und  die  Solvattheorie  der  Losungen.     Zeit.  Elektrochem., 

20,  552  (1914). 
WIOHTMAN,  E.  P.,  WIESEL,  J.  B.,  and  JONES,  H.  C. 

A  Preliminary  Study  of  the  Conductivity  of  Certain  Organic  Acids  hi  Absolute 
Ethyl  Alcohol  at  15°,  25°  and  35°.     Journ.  Amer.  Chem.  Soc.,  36,  2243 
(1914). 
WIOHTMAN,  E.  P.,  DAVIS,  P.  B.,  HOLMES.  A.,  and  JONES.  H.  C. 

Conductibilit6s  et  ViscositSs  des  Solutions  d'lodure  de  Potassium  et  d'lodure 
de  Sodium  dans  lea  Melanges  d'Alcool  Ethylique  et  d'Eau.     Journ.  Chim. 
Phys.,  12,  386  (1914). 
SHAEFFER,  E.  J.,  PAULUS,  M.  G.,  and  JONES,  H.  C. 

Radiometric  Measurements  of  the  lonization  Constants  of  Indicators.     Journ. 

Amer.  Chem.  Soc.,  37,  776  (1915);  Chem.  News,  112,  195  (1915). 
DAVIS,  P.  B.,  and  JONES,  H.  C. 

The  Viscosities  of  Binary  Mixtures  of  the  Associated  Liquids,  Water,  Formio 
Acid  and  Acetic  Acid.     Jour.  Amer.  Chem.  Soc.,  37,  1194  (1915). 


368  THE  NATURE  OF  SOLUTION 

PAULUS,  M.  G.,  HUTCHINSON,  J.  F.,  and  JONES,  H.  C. 

Radiometric  Measurements  of  the  lonization  Constants  of  Indicators. 
Journ.  Amer.  Chem.  Soc.,  37,  1694  (1915). 

DAVIS,  P.  B.,  PUTNAM,  W.  S.,  and  JONES,  H.  C. 

The  Conductivity  and  Viscosity  of  Solutions  of  Electrolytes  in  Formamide. 
Journ.  Franklin  Institute,  567,  Nov.  (1915). 

DAVIS,  P.  B.,  PUTNAM,  W.  S.,  and  JONES,  H.  C. 

Uber  Leitfahigkeit  und  Viskositat  einiger  Rubidium  und  Ammoniumsalze 
in  ternaren  Mischungen  von  Glycerin,  Aceton  und  Wasser.  Zeit.  phys. 
Chem.,  90,  481  (1915). 

WATKINS,  C.  and  JONES,  H.  C. 

Conductivity  and  Dissociation  of  Some  Rather  Unusual  Salts  in  Aqueous 
Solution.  Journ.  Amer.  Chem.  Soc.,  37,  2626  (1915). 

DAVIS,  P.  B.,  and  JONES,  H.  C. 

The  Viscosities  of  Solutions  of  Caesium  Salts  in  Mixed  Solvents.  Journ. 
Amer.  Chem.  Soc.,  37,  2636  (1915). 

HOLMES,  J.  E.  L.,  and  JONES,  H.  C. 

The  Action  of  Salts  with  Water  of  Hydration  and  without  Water  of  Hydra- 
tion  on  the  Velocity  of  Saponification  of  Esters.  Chem.  News,  112,  73 
(1915)  and  Journ.  Amer.  Chem.  Soc.,  38,  105  (1916). 

LLOYD,  H.  H.,  WIESEL,  J.  B.,  and  JONES,  H.  C. 

Conductivities  of  Certain  Organic  Acids  in  Absolute  Ethyl  Alcohol.  Journ. 
Amer.  Chem.  Soc.,  38,  121  (1916). 

PUBLICATIONS  UNDER  THE  AUSPICES  OF  THE  CARNEGIE 
INSTITUTION  OF  WASHINGTON 

JONES,  H.  C.  (with  the  assistance  of  GETMAN,  F.  H.,  BASSE-IT,  H.  P.,  MCMASTER, 
L.,  and  UHLEB,  H.  S.). 

Hydrates  in  Aqueous  Solution.  Carnegie  Institution  of  Washington,  Pub- 
lication No.  60  (1907). 

JONES,  H.  C.  (and  LINDSAY,  C.  F.,  CARBOLL,  C.  G.,  BASSETT,  H.  P.,  BINGHAM, 
E.  C.,  ROUILLEB,  C.  A.,  McMASTEB,  L.,  and  VEAZEY,  W.  R.). 

Conductivity  and  Viscosity  in  Mixed  Solvents.  Carnegie  Institution  at 
Washington,  Publication  No.  80  (1907). 

JONES,  H.  C.,  and  ANDEBSON,  J.  A. 

The  Absorption  Spectra  of  Solutions.  Carnegie  Institution  of  Washington 
Publication  No.  110  (1909). 

JONES,  H.  C.,  and  STBONO,  W.  W. 

A  Study  of  the  Absorption  Spectra  of  Solutions  of  Certain  Salts  of  Potassium, 
Cobalt,  Nickel,  Copper,  Chromium,  Erbium,  Praseodymium,  Neodymium 
and  Uranium  as  affected  by  Chemical  Agents  and  by  Temperature. 
Carnegie  Institution  of  Washington,  Publication  No.  130  (1910). 

JONES,  H.  C.,  and  STBONQ,  W.  W. 

The  Absorption  Spectra  of  Solutions  of  Comparatively  Rare  Salts,  Including 
those  of  Gadolinium,  Dysprosium  and  Samarium.  The  Spectophotography 
of  certain  Chemical  Reactions,  and  the  Effect  of  High  Temperature  on 
the  Absorption  Spectra  of  Non-aqueous  Solutions.  Carnegie  Institution  of 
Washington,  Publication  No.  160  (1911). 


BIBLIOGRAPHY  OF  ARTICLES  AND  BOOKS  369 

JONES,  H.  C.,  and  CLOVER,  A.  MM  HOSFORD,  H.  H.,  HOWARD,  S.  F.,  JACOBSON, 
C.  A.,  KREIDER,  H.  R.,  SHAEFFER,  E.  J.,  SMITH,  L.  D.t  SPRINGER,  A.,  JR., 
WEST,  A.  P.,  WHITE,  G.  F.,  WIGHTMAN,  E.  P.,  and  WINSTON,  L.  G. 

The  Electrical  Conductivity,  Dissociation  and  Temperature  Coefficients  of 
Conductivity  from  0°  to  65°  of  Aqueous  Solutions  of  a  Number  of  Salts 
and  Organic  Acids.  Carnegie  Institution  of  Washington,  Publication 
No.  170  (1912). 

JONES,  H.  C.,  and  STINE,  C.  M.,  PEARCE,  J.  N.,  KREIDER,  H.  R.,  MAHIN,  E.  G., 
SCHMIDT,  M.  R.,  GUT,  J.  S.,  and  DAVIS,  P.  B. 

The  Freezing-Point  Lowering,  Conductivity  and  Viscosity  of  Solutions  of 
Certain  Electrolytes  in  Water,  Methyl  Alcohol,  Ethyl  Alcohol,  Acetone 
and   Glycerol   and    in   Mixtures  of   These  Solvents   with    One   Another. 
Carnegie  Institution  of  Washington,  Publication  No.  180  (1913). 
JONES,  H.  C.,  and  GUT,  J.  S. 

The  Absorption  Spectra  of  Solutions  as  Affected  by  Temperature  and  by 
Dilution.  A  Quantitative  Study  of  Absorption  Spectra  by  Means  of  the 
Radiomicrometer.  Carnegie  Institution  of  Washington,  Publication  No. 
190  (1913). 

JONES,  H.  C.,  WIGHTMAN,  E.  P.,  DAVIS,  P.  B.,  SHAEFFEH,  E.  J.,  HUGHES,  H., 
SMITH,  L.  D.,  HOLMES,  A.,  PAULUS,  M.  G.,  WIESEL,  J.  B.,  and  PUTNAM,  W.  S. 
The  Absorption  Spectra  of  Solutions  as  Studied  by  means  of  the  Radio- 
micrometer;    The    Conductivities,  Dissociations    and    Viscosities    of    So- 
lutions of  Electrolytes  in  Aqueous,   Non-aqueous  and   Mixed  Solvents. 
Carnegie  Institution  of  Washington,  Publication  No.  210  (1915). 
JONES,  H.  C.,  DAVIS,  P.  B.,  SHAEFFER,  E.  J.,  PUTNAM,  W.  S.,  PAULUS,  M.  G., 
LLOTD,  H.  H.,  HOLMES,  J.  E.  L.,  WATKINS,  C.,  WIESEL,  J.  B.,  ORDEMAN,  G.  F., 
CONNOLLT,  G.  C.,  HUTCHINSON,  J.  F.,  and  McCALL,  A.  G. 

Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents,  Radiometric 
Measurements  of  the  lonization  Constants  of  Indicators,  etc.  Carnegie 
Instiuttion  of  Washington,  Publication  No.  230  (1915). 

The  following  papers  will  be  included  in  the  monograph  now  in  prepa- 
ration:— 

DAVIS,  P.  B.,  and  JOHNSON,  H.  I. 

The  Conductivity  and  Viscosity  of  Hydrated  and  Nonhydrated  Salts  in 
Formamide  and  in  Mixed  Solvents  containing  Formamide. 

DAVIS,  P.  B. 

A  Note  on  the  Viscosity  of  Caesium  Salts  in  Glycerol- Water  Mixtures. 
HULBERT,  E.  O.,  and  HUTCHINSON,  J.  F. 

The  Absorption  Coefficient  of  Solutions  for  Monochromatic  Radiation. 
LLOTD,  H.  H.,  and  PARDEE,  A.  M. 

A  Study  of  the  Conductivity  of  Certain  Organic  Salts  in  Absolute  Ethyl 

Alcohol. 
ORDEMAN,  G.  F. 

A  Study  of  the  Dissociating  Powers  of  Free  and  Combined  Water. 

CONNOLLT,  G.  C. 

The  Difference  in  Chemical  Activity  of  Free  and  Combined  Water  as  Illus- 
trated by  the  Effect  of  Neutral  Salts  on  the  Hydrolysis  of  Acetic  Anhydride. 


370  THE  NATURE  OF  SOLUTION 

BOOKS 
JONES,  H.  C. 

The   Freezing-point,   Boiling-point   and   Conductivity   Methods.     Chemical 

Publishing  Company,  Easton  (1897).     Two  editions. 
JONES,  H.  C. 

The  Modern  Theory  of  Solution.    Harper's  Science  Series,  vol.  nr.    American 

Book  Co.  (1899)     Two  editions. 
JONES,  H.  C. 

Practical  Methods  for  Determining  Molecular  Weights.     By  Heinrich  Biltz. 
Translated  by  Jones  and  King.     Chemical  "Publishing  Company,  Easton 
(1899). 
JONES,  H.  C. 

The  Theory  of  Electrolytic  Dissociation  and  Some  of  Its  Applications.     The 

Macmillan  Co.     (1900).     Three  editions. 
JONES,  H.  C. 

Outlines  of  Electrochemistry.     D.  Van  Nostrand  Co.  (1901). 
JONES,  H.  C. 

The  Elements  of  Physical  Chemistry.     Translated  into  Italian  and  Russian. 

The  Macmillan  Co.  (1902).     Four  American  editions. 
JONES,  H.  C. 

Principles  of  Inorganic  Chemistry.     The  Macmillan  Co.  (1903). 
JONES,  H.  C. 

Elements  of  Inorganic  Chemistry.    The  Macmillan  Co.  (1903). 
JONES,  H.  C. 

The  Electrical  Nature  of  Matter  and  Radioactivity.     D.  Van  Nostrand  Co. 

(1906).     Three  editions. 
JONES,  H.  C. 

Introduction   to   Physical    Chemistry.     The    Macmillan    Company    (1910). 

Two  editions. 
JONES,  H.  C. 

A  New  Era  in  Chemistry.     Some  of  the  More  Important  Developments 
in  General  Chemistry  during  the  Last  Quarter  of  a  Century.      D.  Van 
Nostrand  Co.  (1913). 
JONES,  H.  C. 

The  Nature  of  Solution.    D.  Van  Nostrand  Co.  (1917). 


AUTHOR  INDEX 


Alexandrow,  281 

Ambronn,  248 

Anderson,  332,  335,  336,  340,  341,  346 

Archibald,  208 

Arrhenius,  46,  49,  80-84,  86,  87,  89, 

132,  135,  197,  286,  287,  315,  351, 

352 

Aston,  204,  211 
Avogadro,  73,  77-79 

Baker,  9-11 

Barnes,  207 

Bartholi,  81 

Barus,  275 

Bassett,  215,  314 

Baudouin,  257 

Bechhold,  274 

Beckmann,  115,  117,  118,  127,  129, 

130 

Berkeley,  55,  67 
Bernthsen,  150 
Bertheiot,  40-42,  144 
Berthollet,  23,  24 
Berzelius,  27-30,  236 
Beudant,  94,  95 
Biltz,  271,  286 
Bingham,  215,  216 
Bizio,  72 
Blagden,  122 
Blake,  258,  259,  261 
Boerhaave,  21 
Bousfield,  348 
Boyle,  5,  62,  64,  66,  73,  76,  79,  82, 

89,  97,  353 
Braun,  258 
Bredig,  192,  193,  231,  236,  237,  260, 

284,285 
Brown,  R.,  241 
Bruni,  239 
Buchbdck,  348 


Burton,  257,  265 
Biltschli,  274 

Cady,  207 

Cailletet,  3,  278 

Carrara,  203 

Carroll,  214,  215 

Centnerszwer,  206,  207 

Chambers,  306 

Chaperon,  95 

Clausius,  33-35,  49,  81,  83,  164,  165, 

246 

Clover,  316 
Colson,  301,  302 
Coppet,  122 
Cotton,  247,  255 
Coulomb,  209 

Dalton,  25 
Davis,  208,  22S-225 
Davy,  27,  236 
Denison,  348 
DeVries,  67-71,  74 
Dewar,  297 
Dimitrievics,  279 
Dolezalek,  348 
Drude,  210 
Duclaux,  258 
Dutoit,  204,  211 

Ehrenhoft,  243 
Einstein,  243,  246 
Erlich,  286 
Exner,  242,  243 

Faraday,  103,  104, 135, 156-158,  160, 

162,  163,  188,  190,  209,  235 
Favre,  38,  39,  81 
Fick,  92,  93,  97 
Fischer,  288 


372 


AUTHOR  INDEX 


Fitzpatrick,  203 

Fizeau,  248 

Fourcroy,  23 

Fourier,  93 

Franklin,  207 

Frazer,  59,  64,  66,  111,  114,  306 

Freundlich,  265,  266,  274,  276,  280, 

289,293 
Friderich,  204 
Friedmann,  271 

Galvani,  27, 196 

Gay-Lussac,  25,  30,  31,  43,  48,  61, 

62,  64,  66,  72-76,  79,  89,  94-97, 

103,  353 

Getman,  307,  310,  311,  313,  327,  328 
Gibbs,  197 
Gladstone,  39,  281 
Gordon,  287 
Gore,  302 
Gouy,  95 
Graham,  91,  92,  226-232,  234,  240, 

284 

Grotthuss,  25-27,  164 
Guldberg,  35,  36,  48,  72,  135 
Gutbier,  236 
Guthrie,  134 
Gutmann,  10 

Guy,  222-224, 341, 342,  344,  347 
Guye,  211 

Hamburger,  70,  74 
Hantzsch,  150 
Hartley,  55,  67 
Hauer,  303 

Helmholtz,  9,  197,  254 
Henderson,  153 
Henry,  2,  163 
Herzog,  241 
Heydweiller,  174 
Hibbert,  281 
Hippocrates,  4 
Hober,  287 
Hofmeister,  278,  280 
Holland,  59 
Horstmann,  48,  72 
Hosford,  316 
Howard,  316 


Howe,  225 
Hughes,  8,  224 
Humphrey,  120 
Hyde,  207 

Ikeda,  284 

Jacobson,  316 

Jones,  116,  118,  130,  184,  205,  207, 
216,  222-225,  306,  311,  313,  314, 
322,  324,  327,  328,  332,  335,  336, 
338,  340-342,  344,  346,  347 

Kablukoff ,  203,  204 

Kanolt,  348 

Kattein,  238,  279 

Klaproth,  23 

Knight,  306 

Kohlrausch,    40,  81,    172-174,    177, 

178,  181,  195,  219 
Kopp,  35,  93 
Krafft,  290,  292 
Krapiwin,  203,  213 
Kraus,  207 
Krause,  181 
Kreider,  219,  316 
Bolster,  292,  304 

Lagergren,  298 

Landolt,  39 

Lavoisier,  22-24 

Lea,  230,  231,  236 

Le  Blanc,  166,  168,  170 

Le  Chatelier,  298 

L&nery,  5 

Lewis,  208 

Lieben,  95 

Liebig,  93,  105 

Lillie,  238 

Linder,  239,  255,  260-262,  282 

Lindsay,  213 

Lloyd,  209 

Lobry  de  Bruyn,  348 

Lodge,  194-196 

Loeb,  191 

Loew,  230 

Longinescu,  211 

Losev,  293 


AUTHOR  INDEX 


373 


Lottermoser,  273 
Lovelace,  111,  114 
LUdeking,  277,  279 
Ludwig,  96 

Mackay,  184 

Madsen,  286 

Mahin,  220 

Marsden,  302 

Marsh,  9 

Maxwell,  246 

Mclntosh,  208 

McMaster,  216,  328 

Mendel&ff,  44^8,  72,  348,  349 

Michaelis,  272 

Mohler,  120 

Morgan,  348 

Morse,  58-60,  62-64,  66,  97 

Morveau,  22 

Moser,  95 

Mouton,  247,  255 

Murray,  214 

Muthmann,  230 

Myrick,  66 

Nageli,  284 
Neisser,  271 
Nernst,  161-163,  191,  197-199,  209, 

348,  352 
Newton,  20 
Noyes,  A.  A.,  154,  155,  181,  348 

Offer,  134 

Ohm,  93 

Onnes,  175 

Ostwald,  Wilhelm,  105,  129,  130, 
145-150,  159,  161-163,  175,  178- 
180,  186,  187,  201,  315,  317,  352 

Ostwald,  Wolfgang,  233,  280 

Ota,  306 

Paal,289 
Palmaer,  198 
Pappada,  239 
Paracelsus,  5 
Pauchon,  104 
Pauli,  259,  284 
Paulus,  344,  347 


Pearce,  130,  315,  322,  324 

Perrin,  243-246,  254,  257 

Pfeffer,  50-55,  58,  71-74,  77,  97,  238 

Pfund,  297 

Pickering,  348 

Pictet,  3 

Picton,  239,  255,  260-262,  282 

Pincussohn,  272 

Proust,  24 

Putnam,  208,  224 

Quinke,  254,  257 

Ramsay,  119,  120,  211,  212,  214,  243 
Raoult,  43,  44,  49,  81,  106-111,  122- 

127,  132,  133,  238,  240,  281,  311 
Rayleigh,  111 
Raymond,  163,  294 
Reid,  356 
Reinders,  285 
Reinke,  278,  279 
Reuss,  254 
Richartz,  9 
Riesenfeld,  348 
Roberts-Austin,  302 
Robinson,  194 
Rodewald,  238,  279 
Rose,  13,  14 
Rosenstiehl,  72 
Rouiller,  216 
Rowland,  120 
Rudolphi,  180 

Sabanejew,  281 
Sachkanov,  177 
St.  v.  Gorski,  204 
St.  v.  Laszczynski,  204 
Schlamp,  203 
Schmidt,  221,  222,  293 
Selmi,  226 

Shaeffer,  316,  344,  347 
Shields,  211,  212,  214 
Siedentopf ,  247-250 
Smith,  316 
Smits,  290 
Smoluchowski,  243 
Sorensen,  153 
Soret,  75,  76,  96,  97 


374 


AUTHOR  INDEX 


Spring,  300,  301 

Springer,  316 

Stas,  10,  11 

Steele,  348 

Stern,  290 

Stieglitz,  150, 151 

Stokes,  245-247 

Strong,  335,  338,  340,  341,  346,  347 

Svedberg,  237,  242,  243 

Tammann,  104 

Thales,  4 

Thomsen,  41-43,  48,  72 

Thomson,  25,  26,  28,  121,  148,  206, 

209 

Traube,  50 
Tyndall,  247 

Uhler,  330-332,  346 

Valentine,  4 

Valson,  36,  38,  39,  44,  49,  81 

Van  Bemmelen,  276,  298 

Van  der  Waals,  89,  353 

Van  Helmont,  4,  5 

Van't  Hoff,  31,  48,  63,  72-74,  76-80, 
82,  84,  86,  89,  118,  128,  129,  132, 
135,  181,  196,  197,  298-304 

Vaubel,  298 

Veazey,  217,  222,  224,  225 

VioUe,  302 

Vollmer,  203 

Volta,  27,  196 


Von  Berneck,  284,  285 
Von  Meyer,  373 
Von  Schroder,  277,  278 
Von  Weimarn,  283 

Waage,  35,  36,  48,  72,  135 
Walden,  205,  206 
Walker,  105 
Wallerius,  22 
Wanklyn,  8 
Washburn,  348 
Wenzel,  280 
Werner,  204 
West,  316 

Whitney,  258,  259,  261 
Wiedemann,  254,  279 
Wiener,  242 
Wiesel,  209 
Wight,  316 
Wightman,  209,  316 
Williamson,  31-33,  48,  81 
Winston,  316 
Witt,  292 
Wladimiroff,  70 
Wohler,  93 
Wroblewski,  3 
WiOlner,  95,  104 

Zanniovich-Tessarin,  205 
Zelinsky,  203,  213 

Zsigmondy,  236,  247,  248,  250-252, 
273 


SUBJECT  INDEX 


ABNORMAL  electrolytes,  206 
Absorption  of  light  by  solutions,  330 

—  spectra   of    solutions,    effect    of 
dehydrating  agents  on,    332 

effect  of  solvent,  332,  334 

of  neodymium  salts,  333, 338 

of  uranium  salts,  335 

relative  transparency  to  sol- 
vent, 342 

solvent  bands,  332,  335,  339 

use  of  radiomicrometer,  342 

Acetone,  conductivities  in  mixtures 

containing,  215,  224 
Acid,  definition  of,  136 
Acids,  dissociation  of,  183 
Active  and  inactive  molecules,  83 
Activity  coefficients   (Arrhenius),  83 
Adsorption,  application  to  high  vacua, 
297 

—  effect  of  solvent  on,  295 
temperature  on,  296 

—  history  of,  294 

—  in  soils,  288 

—  theories  of,  297 

—  theory  of  colloids,  282 
Aggregation,  states  of,  2 

Alcohol,   conductivities   in  mixtures 

containing,  213,  215,  221 
"Alkahest,"  5 
Ammonia,     dissociating     power     of 

liquid,  207 
Amphoteric  compounds,  and  biology, 

152 

dissociation  of,  151 

Amplitude  of  Brownian  movement, 

242 
Analogy  between  solutions  and  gases, 

30,  42,  72 

Antitoxins  and  toxins,  286 
Association,  and  dissociating  power 

of  liquids,  211 


Association,  effect  of  temperature  on, 
212 

—  factors  of  liquids,  212 

Atomic  volume,  and  viscosity,  218 

—  mass  from  Brownian  movement, 
246 

BASE,  definition  of,  137 
Bases,  dissociation  of,  183 
Berthollet's  chemical  theory  of  solu- 
tion, 23 

Berzelius'  electrochemical  theory,  27 
Boiling-point  apparatus,  (Jones),  116 

—  method  (Beckmann),  115 

electrolytic    dissociation    by, 

118 

molecular  weights  by,  117 

Brownian  movement,  241 

amplitude  of  vibrations  in,  242 

and  kinetic  theory,  246 

and  precipitation  of  colloids, 

262 

and  ultramicroscopy,  251 

and  viscosity  of  media  (Sved- 

berg),  242 

mass  of  atom  from,  246 

Perrin's  work  on,  243 

theories  of,  243 

velocity  of  particles  in,  242 

CAPILLARY  moduli,  law  of  (Valson), 

36 

Carbon  dioxide,  geological  action  of,  13 
Catalysts,  action    of  "poisons"  on, 

285 

Cataphoresis,  254 
Cells,  primary,  196 
Chemical  action,  gravitational  theory 

of  (Newton),  20 
Chemical  equilibrium,  importance  of 

recognized  (Williamson),  31 


376 


SUBJECT  INDEX 


Chemical    inactivity  in   absence   of 

water,  6 
of  hydrochloric  acid  and 

carbonates,  8 

of  acids  on  litmus,  9 

of  hydrochloric  acid  and 

ammonia,  9 
of    sulphuric    acid    and 

sodium,  11 

Colloidal  particles,  size  of  and  ultra- 
microscope,  250 
Colloidal      preparations,      chemical 

methods,  234 
electrical  methods,   234,   236, 

284 

hydrolysis  methods,  234 

reduction  methods,  235 

Colloidal  solutions,  as  catalysts,  285 

fundamental  properties  of,  237 

of  metals,  231,  236,  284 

precipitation  of,  260 

action   of  electrolytes,  261, 

264 
and  valency  of  precipitating 

ion,  265,  268 

microscopic  phenomena,  262 

preparation  of,  234 

Colloids,  adsorption  theory  of,  282 

—  and  agriculture,  288 

—  and  mineralogy,  289 

—  and  pharmacology,  287 

—  early  history  of,  226 

—  electrical  properties  in  nonaque- 
ous  solutions,  257 

—  Graham's  work  on,  227 

—  industrial  applications  of,  289 

—  mutual  precipitation  of,  271 

—  relation    to    physiological    chem- 
istry, 284 

—  suspension  theory  of,  282 

—  theory  of  precipitation  of,  263,  283 
Color  in  solution,  cause  of,  145 

and  electron  theory,  147 

and   theory   of  electrolytic 

dissociation,  146 

Conductivities,  in  acetone  mixtures, 
215 

—  in  alcohol  mixtures,  213 


Conductivities,  in  glycerol,  221 

—  in  mixed  solvents,  213 
Conductivity,     electrical,     at     high 

temperatures,  181 

theory  of,  171 

method,  Kohlrausch,  40,  172 

of  solutions,  law  of,  177 

—  of  suspensions,  258 

—  water,  174 
Cryohydrates,  133 

DECOMPOSITION    values   of   electro- 
lytes, 167 
Dialysis,  231 

—  Graham's  method,  228 
Dielectric    constants,    determination 

of  (Drude),  210 
effect  of  on  dissociating  power 

of  liquids,  209 

of  liquids,  211 

Diffusion,  90 

—  and  osmotic  pressure,  96 

—  effect  of  mass  on,  95 

—  generalization  of  Fick,  93 

—  Graham's  work  on,  91,  227 

—  in  solids,  301 

—  of  sols,  240 

—  temperature  coefficients  of,  96 
Dilution  law,  Ostwald,  179 

Rudolphi,  180 

Dissociating  power,  and  association, 

211 

and  dielectric  constants,  209 

of  liquid  ammonia,  207 

of  liquids,  relative,  203 

Walden's  work  on,  205 

Dissociation,  electrolytic,  and  dilu- 
tion, 179 

by  boiling-point  method,  118 

comparison    of    freezing-point 

and    conductivity    measurements 
of,  315 

by  freezing-point  method,  129 

in  acetone,  204 

in  alcohols,  203,  209 

in  ether,  204 

in  fonnamide,  208,  224 

in  formic  acid,  205 


SUBJECT  INDEX 


377 


Dissociation,  electrolytic,  in  hydro- 
cyanic acid,  206 

in  hydrogen  dioxide,  207 

in  inorganic  liquids,  205 

in  nonaqueous  solvents,  219 

in  pyradine,  204 

in  ternary  mixtures,  220 

in  water,  183 

of  acids,  183 

of  bases,  183 

of  complexes,  184 

of  salts,  184 

theory  of,  82 

—  hydrolytic,  153 
Dyeing,  291 

—  and  adsorption,  293 

—  and  chemical  action,  292 

—  and  solid  solutions,  293 
Dynamics,  chemical,  importance  rec- 
ognized, 143 

ELECTRICAL    charge    and    precipita- 
tion of  colloids,  271 

—  dispersion,  236 

—  properties  of  sols,  252 
Electrochemical  theory  of  Berzelius, 

27 
Electroendosmose  of  sols,  253 

—  effect  of  electrolytes  on,  254 
Electrolysis  and  conductance,  161 

—  primary  of  water,  170 

—  theories  of,  27,  33,  164,  166 

—  theory  of  and  decomposition  val- 
ues, 167 
Electrolytes,  abnormal,  206 

—  action  of  on  colloids,  261 

—  decomposition  values  of,  167 

—  osmotic  pressure  of,  64 
Electrolytic  dissociation,  by  boiling- 
point  method,  118 

by  conductivity  method,  178 

by  freezing-point  method,  129 

early   conceptions   of    (Grott- 

huss),  25 

methods  of  measuring,  83 

—  theory  of,  82,  87,  351 
Electrostenolysis,  258 
Emulsions,  233 


Emulsions,  electrical  properties  of,  259 

—  precipitation  of,  273 
Emulsoids,  233 

—  osmotic  pressure  of,  238 

—  precipitation  of,  270,  273 
Enzymes,  285 

Ether  formation,  theory  of  (William- 
son), 31 

FARADAY'S  laws,  156 

Ferments,  organic  and  inorganic,  284 

—  catalytic  action  of,  285 
Formamide,  dissociation  hi,  208,  224 
Formic  acid,  dissociation  in,  205 
Freezing-point    investigations,    122, 

130 

—  lowering,  abnormal,  306 

and    water   of   crystallization, 

308 

bearing  on  hydration,  309 

by  sols,  239 

in  solid  solutions,  304 

Raoult's  law  of,  43,  126 

—  method  (Beckmann),  127 

molecular  weights  by  128 

Fusible  metals,  304 

GASES,    analogy    between    solutions 
and,  30,  42,  72 

—  liquefaction  of,  3 
Gelatine,  270 
Gelation,  234,  270 
Gels,  232,  274 

—  honeycomb  structure  of,  275 

—  imbibition  by,  277 

—  less  elastic,  276 

—  more  elastic,  277 

—  physical  properties  of,  275 

—  solventation  of,  276 
Glycerol,  conductivities  in,  221 

—  viscosities  in,  222 
Gold  number,  273 

HEAT  of  neutralization,  139 
Hydrates,  composition  of,  310 

—  early  views  on,  41 

—  of  sulphuric  acid,  45,  47,  349 

—  theories  of,  44,  307,  348 


378 


SUBJECT  INDEX 


Hydration,  and  atomic  volume,  322 

—  and  ionic  velocity,  193,  324 

—  and  osmotic  pressure,  62 

—  and   temperature    coefficients   of 
conductivity,  316 

—  of  gels,  276 

—  relative  of  anion  and  cation,  324 
Hydrolysis,  153 

Hydrolytic  dissociation,  153 

effect  of  temperature  on,  153 

Hydrosols  and  hydrogels,  232 

IMBIBITION,  277 

—  calorimetry  of,  279 

—  effect  of  solutions  on,  280 

—  pressure  from,  278 

—  rate  of  by  gels,  278 

—  volume  changes  in,  277 
Indicators,   ionization   constants  of, 

151 

—  Ostwald  theory  of,  149 

—  Stieglitz's  work  on,  150 
Ionic  theory,  87 

Ions,  83 

—  absolute  velocity  of,  194 

—  apparatus  for  measuring  velocity 
of,  188 

—  effect  of  concentration  and  tem- 
perature on  velocity  of,  190 

—  hydration  of  and  atomic  volume, 
322 

—  modes  of  formation,  186 

—  relative  velocities  of,  188,  192 

—  velocity  of  and  hydration,  193,  324 
Ionization  constants,  84 

—  of  indicators,  151 
Isohydric  solutions,  185 
Isosmotic  solutions,  69 

Isotonic  coefficients  (De  Vries),  69 

KINETIC  theory,  and  Brownian  move- 
ment, 246 
Kohlrausch  conductivity  method,  172 

—  law,  40,  177 

LAVOISIER'S  conceptions  of  solution, 

22 
Liquefaction  of  gases,  3 


Liquids,  association  factors  of,  212 

—  dielectric  constants  of,  211 

—  relative  dissociating  power  of,  203 

—  solvent  power  of,  202 

—  Walden's    work    on    dissociating 
power  of,  205 

Lyophile  —  lyophobe  sols,  234 

MASs-action,  law  of  (Guldberg  and 

Waage),  14,  36 
Mercury,  molecular  weights  of  metals 

in,  119 
Metals,   colloidal  solutions  of,  231, 

236,  284 

—  molecular  weights  in  mercury,  119 

—  solution  tension  of,  197,  200 
Migration  velocities  of  ions,  40,  177, 

188,  192,  195 

Mixed  solvents,  conductivities  in,  213 
Moduli  of  coersion  (Farve  and  Val- 

son),  38 
Molecular  weights,  by  boiling-point 

method,  117   • 

by  freezing-point  method,  128 

of  metals  in  mercury,  119 

Mordaunts,  292 

NEGATIVE  sols,  255 
precipitation  of,  266 

—  viscosity  coefficients,  218,  223 
Neutralization,  137 

—  heat  of,  139 

Nomenclature  of  colloidal  chemistry, 
231,  233 

—  of  ultramicroscopy,  249 
Nonaqueous  solvents,  dissociation  in, 

202,  204,  219 

OSMOTIC  pressure,  and  diffusion,  96 

and  Boyle's  law,  73 

and  Gay  Lussac's  law,  74 

application  of  gas  laws  to,  72 

cell  for  high  (Frazer),  66 

durability  of  cells,  66 

in  solid  solutions,  300 

measurements  of  Pfeffer,  50 

membranes  by  electrical  endos- 

mose  (Morse),  58 


SUBJECT  INDEX 


379 


Osmotic  pressure,  Morse's  method  of 

measuring,  59 

of  cane  sugar,  60 

of      concentrated      solutions, 

(Berkley  and  Hart),  55 

of  electrolytes,  64 

of  emulsoids  and  suspensoids, 

238,  239 

of  glucose,  63 

relation  to  gas  pressure,  60,  77 

relative,  67,  70 

r61e  in  different  sciences,  99 

Traube's  method  of  measuring, 

50 
Ostwald  dilution  law,  179 

—  theory  of  indicators,  149 

POISONS,  catalytic,  285 

Positive  sols,  255 

precipitation  of,  265 

Precipitation  of  colloids,  and  Brown- 
ian  movement,  262 

protective  effect,  273 

valency  rule,  268 

Primary  cells,  196 

Nernst's  work  on  198 

Protective  effect  and  colloidal  pre- 
cipitation, 269,  273 

Purple  of  Cassius,  236 

RADIOMICROMETER,  342 
Raoult's  laws,  111,  127 
Resonance,  145,  330 

SALTS,  dissociation  of,  184 

—  thermoneutrality  of,  142 
Selective    absorption    of    light    by 

solutions,  330 
Semicolloids,  289 
Semipermeable  membranes  (Morse), 

58 

preparation  of  (Traube),  50 

Silver,  colloidal,  (Lea),  230 

—  ion  and  theory  of  hydration,  326 
Soaps,  290 

Size  of  colloidal  particles  and  ultra- 
microscopy,  250 
Solation,  234 


Solids,  diffusion  in,  302 

—  osmotic  pressure  in  mixtures  of, 
301 

—  vapor-tension,  lowering  of,  303 
Solid  solutions,  kinds  of,  299 

Sols,  boiling-point,  freezing-point, 
and  vapor  pressure  of,  239 

—  diffusion  of,  240 

—  electrical  properties,  252 

—  mutual  precipitation,  271 

—  precipitation  of,  260 

Solution,  Lavoisier's  conception  of,  22 

—  law  of  electrical  conductivity  of, 
177 

—  modern  definition  of,  2,  6 

—  r61e  of  in  different  sciences,  12-16 

—  theories  of,  chemical,  Klaproth,  23 
BerthoUet,  23 

—  Kopp,  35 

Thomson,  25 

mechanical,  Lavoisier,  23 

Fourcroy,  23 

17th  and  18th  century,  Newton, 

20 

Boerhaave,  21 

Wallerius,  22 

Solutions,  and  gases,  30,  42,  72,  78 

—  as  compounds  (Kopp),  35 

—  fundamental  properties,  132 

—  isohydric,  185 

—  law  of  electrical  conductivity  of, 
177 

Solution-tension  of  metals,  197 

proof  of,  198 

values  for,  199 

Solvate  theory,  and  chemistry  in 
general,  355 

relation  to  theory  of  electro- 
lytic dissociation,  357 

Solvation,  and  selective  absorption  of 
light,  331 

—  in  nonaqueous  solutions,  329 

—  in  the  alcohols,  327 

—  summary  of  evidence  for,  345 
Solvent  bands,  332,  335,  337,  341 
Solvents,  nonaqueous,  dissociation  in, 

202,  219 
Soret,  principle  of,  75,  96 


380 


SUBJECT  INDEX 


Surface  tension,  263,  285,  297 
Suspensions,  233 

—  action  of  electrolytes  on,  261,  269 

—  conductivity  of,  258 
Suspension  theory  of  colloids,  282 
Suspensoids,  233 

—  osmotic  pressure  of,  239 

—  precipitation  of,  260 

TANNING,  291 

Temperature  coefficients  of  con- 
ductivity, and  relative  hydrating 
power,  321 

of  strongly  hydrated  salts, 

319 

of  weakly  hydrated  salts, 

318 

—  effect  of  on  adsorption,  296 
on  association,  212 

Ternary  mixtures,  dissociation  in, 
220 

Theory  of  electrolytic  dissociation, 
82,  87,  351 

Thennochemical  investigations,  Ber- 
thollet,  40 

Thomson,  42 

Thermoneutrality  of  salts,  142 

Toxins  and  antitoxins,  286 

Transmission  of  light,  relative  by  sol- 
vent and  solution,  345 

Tyndall  effect,  247 

ULTRAMICROSCOPB,  247 

—  applications  of,  249 
Ultramicroscopy,    nomenclature    of, 

249 
Universal  solvent,  5 

VACUA,  high  by  adsorption,  297 
Valence,  nature  of  chemical,  158 
Vapor   pressure    lowering,    in    solid 
solutions,  303 


Vapor-tension,  early   work  on,  103- 
105 

effect  of  concentration  on,  107 

solute,  109 

solvent,  110 

temperature,  108 

lowering  of  by  sols,  239 

measurements,  Walker,  105 

Frazer-Lovelace  method,  111 

of  nonaqueous  solutions,  106 

Raoult's  work  on,  106 

Velocity  of  ions,  188 

—  absolute,  194 

—  apparatus  for  measuring,  (Jones), 
189 

—  effect  of  concentration  and  tem- 
perature on,  190 

—  relative,  188,  192 

and  atomic  volume,  193 

Velocity  of  particles  in  cataphoresis, 

255 
Viscosity,  217 

—  and  Brownian  movement,  242 

—  negative  coefficients  of,  218,  223 

—  of  glycerol  solutions,  222 

—  of  rubidium  and  caesium  salts,  223 

—  relation  to  atomic  volume,  218 

WATER,  as  the  original  element,  4 

—  chemical  and  physical  properties 
of,  17 

—  chemical    inactivity,   in    absence 
of,  7 

—  dissociation  of  solutions  in,  183 

—  importance  of  in  chemical  reac- 
tions, 138 

—  of  crystallization,  and  hydration, 
308 

and  temperature,  314 

—  purification  of  for  conductivity,  174 

—  theory  of  primary  electrolysis  of, 
170 


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HENEW  -VL  i£S  WAY  BE  MADE  4  DAYS  PRIOR 

LOAN  PERIODS  ARE  1 -MONTH.  3-MONTHS.  AND  1-YEAtt 
RENEWALS.  CALL  <415)  642-^406 


TO  DUE  DATE. 


DUE  AS  STAMPED  BELOW 

JAN  03  1991 

MO  DISC  WV  ? 

131 

UNIVERSITY  OF  CALIFORNIA,  BERKELEY 
FORM  NO.  DD6,  60m,  1/83          BERKELEY,  CA  94720 


•37613,3 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


